Covalent Bonding

1. Covalent Bonding • In nature, only the noble gas elements exist as uncombined atoms. • They are monoatomic - consist of single atoms. • The octet rule states that chemical compounds form so each atom (through gaining, losing, or sharing electrons) will have 8 valence electrons • Exception: atoms trying to be like helium

2. Covalent Bonding • Some elements will lose or gain electrons • To form ionic compounds • Some elements will share electrons • To form molecular compounds

3. Covalent Bonding • A covalent bond is the sharing of electrons between two atoms • A neutral group of atoms held together by covalent bonds is called a molecule

4. Covalent Bonding • There are some elements in the periodic table that do not exist alone • Diatomic elements – two atoms of the same element covalently bonded together • The 7 naturally occurring diatomics are: • H2, N2, O2, F2, Cl2, Br2, I2

5. Covalent Bonding • Molecules can be made of the same element or different elements • A compound composed of molecules is called a molecular compound. • All molecular compounds have a molecular formula.

6. Covalent Bonding • A molecular formula reflects the actual number of atoms in each molecule. • The subscripts are not necessarily the lowest whole-number ratios. • For example, the formula for peroxide is H2O2 • Each molecule of peroxide contains 2 hydrogen atoms and 2 oxygen atoms

7. Covalent Bonding • We use Lewis Dot Diagrams to show covalent bonding • However, we do not need to put the dots in the same order as before • We need to put them in singles before we can pair them up

8. Covalent Bonding • In the F2 molecule, each fluorine atom contributes one electron to complete the octet. • Notice that the two fluorine atoms share only one pair of valence electrons. That is a single covalent bond • When we show the bonding, we use Lewis structures • Structural formulas are a neater way to show bonding

9. Covalent Bonding • A pair of valence electrons that is not shared between atoms is called an unshared pair • In F2, each fluorine atom has three unshared pairs of electrons.

10. Covalent Bonding • A double covalent bond is a bond that involves two shared pairs of electrons.

11. Covalent Bonding • Similarly, a bond formed by sharing three pairs of electrons is a triple covalent bond.

12. Methane (structural formula) Methane (Lewis Structure) Covalent Bonding • Electron dot structures fail to reflect the three-dimensional shapes of molecules. • The electron dot structure and structural formula of methane (CH4) show the molecule as if it were flat and merely two-dimensional.

13. Covalent Bonding • To determine the 3D shape of the molecule, we use VSEPR (valence shell electron pair repulsion) theory • The theory states that repulsion between sets of valence level electrons surrounding an atom causes these sets to be oriented as far apart as possible • Or simply, unshared pairs of electrons want to be as far apart as possible

14. Covalent Bonding • The hydrogens in a methane molecule are at the four corners of a geometric solid called a regular tetrahedron. • In this arrangement, all of the H–C–H angles are 109.5°, the tetrahedral angle.

15. Covalent Bonding Unshared electron pair • The molecule ammonia (NH3) is trigonal pyramidal shape. • However, one of the valence-electron pairs is an unshared pair and itrepels the bonding pairs, pushing them together. • The measured H—N—H bond angle is only 107°, rather than the tetrahedral angle of 109.5°. 107°

16. Covalent Bonding • The water molecule is planar (flat) but bent. • With two unshared pairs repelling the bonding pairs, the H—O—H bond angle is compressed to about 105°. 105°

17. Carbon dioxide (CO2) 180° No unshared electron pairs on carbon Covalent Bonding • CO2 is a linear molecule • The carbon in a carbon dioxide molecule has no unshared electron pairs. • The double bonds joining the oxygens to the carbon are farthest apart when the O=C=O bond angle is 180°

18. Covalent Bonding • Even when electrons are being shared, the sharing is not always equal

19. Covalent Bonding • Equal pull between atoms results in a nonpolar covalent bond • Unequal pull between atoms results in a polar covalent bond

20. Covalent Bonding • Equal sharing results in no partial charges on atoms • The unequal sharing of electrons results in one atom having a partial negative and the other atom having a partial positive charge H—H δ+ δ– H—O

21. Covalent Bonding • The electronegativity difference between two atoms tells you what kind of bond is likely to form.

22. H—O Covalent Bonding • The polar nature of the bond may also be represented by an arrow pointing to the more electronegative atom.

23. Covalent Bonding • A molecule that has polar bonds can be overall a nonpolar molecule • Some shapes that can yield nonpolar molecules are: • Linear • Trigonal planar • Tetrahedral

24. Ionic Bonding • An ionic compound is a compound composed bonds between cations and anions. • Anions and cations have opposite charges and attract one another by means of electrostatic forces. • The electrostatic forces that hold ions together in ionic compounds are called ionic bonds.

25. Bonding This table summarizes some of the characteristic differences between ionic and molecular substances.

26. Bonding Collection of water molecules Array of sodium ions and chloride ions Formula unit of sodium chloride Molecule of water Chemical Formula Chemical Formula H2O NaCl

27. Covalent Bonding • The forces between atoms within a molecule are called intramolecular forces • The forces between neighboring molecules are called intermolecular forces

28. Covalent Bonding • There are three types of intermolecular forces • Dipole • Hydrogen bonding • Dispersion forces • These forces are weaker than intramolecular forces, but can have a significant affect a substance’s properties

29. Covalent Bonding • Dipole interactions occur when polar molecules are attracted to one another. • The electrical attraction occurs between the oppositely charged regions of polar molecules. • Dipole interactions are similar to, but much weaker than, ionic bonds.

30. Covalent Bonding • Hydrogen bonding occurs when hydrogen is covalently bonded to O, N, or F and also attracted to an unshared electron pair on an O, N, or F from a neighboring molecule • Hydrogen bonds are the strongest of the intermolecular forces.

31. Covalent Bonding • Dispersion forces are caused by the motion of electrons. • They occur in all molecules, even between nonpolar molecules. • When the moving electrons happen to be momentarily more on the side of a molecule closest to a neighboring molecule, their electric force influences the neighboring molecule’s electrons to be momentarily more on the opposite side. • The strength of dispersion forces generally increases as the number of electrons in a molecule increases.

32. Covalent Bonding • Fluorine and chlorine have relatively few electrons and are gases at ordinary room temperature and pressure because of their especially weak dispersion forces. • Bromine molecules therefore attract each other sufficiently to make bromine a liquid under ordinary room temperature and pressure. • Iodine, with a still larger number of electrons, is a solid at ordinary room temperature and pressure.