Covalent Bonding

1. Covalent Bonding Lesson 2

2. Covalent Bonds • Atoms can share electrons to attain the electron configuration of noble gases (8 valence electrons). • Shared electrons result in covalent bonds. • Covalent bonds usually occur between elements that are relatively close to each other on the Periodic Table. • The majority of covalent bonds form between nonmetals. • A molecule forms when two or more atoms bond covalently.

3. Formation of a Covalent Bond • Covalent bonds form when the attraction between the positive nucleus of one atom and the electrons of another atom are greater than the repulsive force between electrons of both atoms. • Many elements form diatomic molecules that consist of two atoms of the same element. • The atoms share the necessary electrons to form a covalent bond.

4. Diatomic Molecules • Oxygen – O2 • Hydrogen – H2 • Nitrogen – N2 • Fluorine – F2 • Chlorine – Cl2 • Bromine – Br2 • Iodine – I2

5. Single Covalent Bonds • A single covalent bond forms when a single pair of electrons is shared by the atoms. • The shared electron can be depicted in electron dot structure as two dots or as a line. • Lewis structures use electron dot configuration to show how electrons are arranged in molecules.

6. 7A elements share one electron to form one covalent bond. • 6A elements share two electrons to form two covalent bonds. • 5A elements share three electrons to form three covalent bonds. • 4A elements share four electrons to form four covalent bonds.

7. Sigma Bonds • Single covalent bonds are also called sigma bonds when the electrons being shared are in an area centered between the two atoms. • Methane (CH4), ammonia (NH3), and water (H2O) all form sigma bonds.

8. Multiple Covalent Bonds • When atoms share more than one pair of electrons between two atoms, they form multiple covalent bonds. • Double covalent bonds occur when two pairs of electrons are shared. • Triple covalent bonds occur when three pairs of electrons are shared.

9. Pi Bonds • Pi bonds form when parallel orbitals overlap to share electrons. • Double covalent bonds consist of one sigma bond and one pi bond. • Triple covalent bonds consist of one sigma bond and two pi bonds. • A pi bond always accompanies a sigma bond when forming a double or triple bond.

10. Strength of Covalent Bonds • Depends on the distance between bonded nuclei. • Bond length is the distance between bonded nuclei and this length decreases as more electrons are shared. • The shorter the bond length, the stronger the bond. • Energy must be used to break bonds and the energy necessary is called the bond dissociation energy.

11. Naming Covalent Molecules • Binary compounds • The first element in the formula is always named first, using the entire element name. • The second element in the formula is named using the root of the element and adding the suffix –ide. • Prefixes are used to indicate the number of atoms of each type that are present in the compound. • Don’t use mono- with the first element name. • Avoid awkward pronunciations.

12. Common Prefixes • Mono- one • Di- two • Tri- three • Tetra- four • Penta- five • Hexa- six

13. Examples • CCl4 - Carbon tetrachloride • As2O3 – Diarsenic trioxide • CO – Carbon monoxide • SO2 – Sulfur dioxide • NF3 – Nitrogen trifluoride

14. Do Covalent Molecule Problems

15. Naming Acids • Hydrogen compounds that are dissolved in water are acids. • Binary acids are composed to hydrogen and another element. • Use the prefix hydro- for hydrogen. • Use the root of the second element plus the suffix –ic, followed by the word acid. • Example: HF is hydrofluoric acid

16. Naming Polyatomic Acids • Use the prefix hydro-for hydrogen. • Use the root of the polyatomic ion name and the suffix –ic. • Add the word acid. • Example: HCN is hydrocyanic acid.

17. Naming Oxyacids • Oxyacids contain oxyanions. • Identify the oxyanion. Use the root of the oxyanion, a suffix, and the word acid. • -ate is replaced with –ic. • -iteis replaced with –ous. • Examples: • HNO2: nitrous acid • HNO3: nitric acid

18. Practice Problems • Name the following acids: • HI • HClO3 • HClO2 • H2SO4 • H2S

19. Writing Names from Formulas • Look at the formula of the molecule. • Does the compound form an acidic aqueous solution? • Acid: name as an acid • Not an acid: name the first element in the molecule. Use a prefix if the number of atoms is greater than one. Name the second element using a prefix, the root of the element, and –ide.

20. Does the acid have oxygen in the formula? • No oxygen: hydro- plus the root of the second element plus –ic, then acid. • Oxyanion: root of the oxyanion plus –ic, if the anion ends in –ate, or plus – ous, if the anion ends in –ite, then acid. • Examples: • HClO3 = Chloric acid • H2SO4 = Sulfuric acid • HNO2 = Nitrous acid

21. Writing Formulas from Names • Look at the name • If it has an acid in the name, then it has hydrogen as the first element. • If it has a polyatomic name, then look at the suffix –ous or –ic to determine whether it is an –ite or -ate. • Look at prefixes, use the correct subscript to indicate how many of each atom is present.

22. Practice Problems • Write the molecular formula for each listed compound. • Dinitrogen trioxide • Nitrogen monoxide • Hydrochloric acid • Chloric acid • Sulfuric acid • Sulfurous acid

23. Structural Formulas • Structural formulas use letter symbols and bonds to show relative positions of atoms. • Lewis structures can be determined using the following steps: • Predict the location of certain atoms. • Hydrogen is always a terminal, or end, atom. • The atom with the least attraction for shared electrons in the molecule is the central atom.

24. Find the total number of electrons available for bonding. This total is the number of valence electrons in the atoms in the molecule. • Determine the number of bonding pairs by dividing the number of electrons by two. • Place one bonding pair between the central atom and each of the terminal atoms. • Continue adding bonding pairs to the atoms to complete the octet rule for each atom.

25. Practice Problems • Draw a Lewis structure for each of the following molecules: • NF3 • CS2 • BH3

26. Resonance Structures • Resonance occurs when more than one valid Lewis structure can be written for a molecule or ion. • Resonance structures differ only in the position of the electron pairs, never the atom positions. • The location of the lone pairs and bonding pairs differs in resonance structures.

27. Exceptions to the Octet Rule • Odd number of valence electrons cannot form an octet around each atom. • Some compounds form with fewer than eight electrons present around an atom. • Some compounds have central atoms that have more than eight valence electrons because of the d orbitals.

28. Molecular Shape • VSEPR: Valence Shell Electron Pair Repulsion model is based on an arrangement that minimizes the repulsion of shared and unshared pairs of electrons around the central atom. • The repulsions result in atoms existing at fixed angles to each other. • The angle formed by any two terminal atoms and the central atom is a bond angle. • Shared electron pairs repel each other. • Lone electrons occupy a slightly larger orbital than shared electrons and push the shared electrons slightly together.

29. Common Molecular Shapes • Linear: 2 pairs of shared electrons = 180o • Trigonal: 3 pairs of shared electrons = 120o • Tetrahedral: 4 pairs of shared electrons = 109.5o • Trigonal pyramidal: 3 pairs of shared electrons and 1 pair of lone electrons = 107.3o • Bent: 2 pairs of shared electrons and 2 pairs of lone electrons = 104.5o • Trigonalbipyramidal: 5 pairs of shared electrons = 90oand 120o • Octahedral: 6 pairs of shared electrons = 90o

30. Hybridization • Process by which atomic orbitals are mixed to form new, identical hybrid orbitals. • The s and p orbitals are mixed and electrons from each orbital is shared. • Each hybrid orbital contains one electron that it can share with another atom.

31. Electronegativity and Polarity • Electronegativity indicates the relative ability of an atom to attract electrons in a chemical bond. • Covalent bonds between identical atoms are non-polar because there is no difference in electronegativity. • Covalent bonds between different atoms are polar because there is a difference in electronegativity.

32. Polar Covalent Bonds • The more electronegative atom is located at the partially negative end, while the less electronegative atom is found at the partially positive end. • This polar bond is a dipole. • Non-symmetric covalent molecules are polar while symmetric covalent molecules are not. • Polar substance can dissolve only in other polar substances.

33. Properties of Covalent Compounds • Relatively low melting and boiling points • Relatively soft • Covalent network solids are typically brittle, nonconductors of electricity or heat, and extremely hard.