1. Bonding IB Chemistry 2 Robinson High School Andrea Carver

2. Chemical Bonding Elements readily combine with one another to form compounds. A Chemical Bond is the attractive force which holds these elements together in a compound. Compounds will have different characteristics than their component elements. Types of Bonding: Ionic, Covalent, Metallic

3. Ionic Bonding: IB Objectives 4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions. 4.1.2 Describe how ions can be formed as a result of electron transfer. 4.1.3 Deduce which ions will be formed when elements in Groups 1,2, and 3 lose electrons. 4.1.4 Deduce which ions will be formed when elements in Groups 5, 6, and 7 gain electrons. 4.1.5 State that transition elements can form more than one ion. 4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the Periodic Table or from their electronegativity values. 4.1.7 State the formula of common polyatomic ions formed by non- metals in Periods 2 and 3. 4.1.8 Describe the lattice structure of ionic compounds.

4. Ion Formation Ions form when electrons are lost or gained. Typically, atoms are neutral. An ion is an atom that carries an electric charge. ◦ Cation-positively charged, lost electrons ◦ Anion- negatively charged, gained electrons

5. Ion Formation The periodic table may be used to predict the type/charge of ions formed by an element. ◦ Electrons will be gained/lost so that the atom obtains 8 valence electrons. ◦ Elements with less than 4 valence electrons will lose electrons to become cations. ◦ Elements with more than 4 valence electrons will gain electrons to become anions.

6. Ion Formation Transition Elements- These elements are able to gain or lose a variable number of electrons due to their electron configurations. ◦ Examples:  Fe: Fe2+ and Fe 3+  Cu: Cu2+ and Cu3+

7. Ion Formation Polyatomic Ions- Compounds which exist as ions. ◦ Examples:  Nitrate  Hydroxide  Hydrogencarbonate  Carbonate  Sulfate  Phosphate  Ammonium

8. Ionic Bond Formation An Ionic Bond is an electrostatic attraction between oppositely charged ions. An Ionic Compound is held together by this type of force. ◦ An ionic compound will be electrically neutral because the charges of the ions involved will be balanced.

9. Ionic Character The Ionic Character of the compound may be predicted by: ◦ Position of the elements on the periodic table:  Metals + Nonmetals  Bottom Left + Top Right ◦ Electronegativity:  Difference in electronegativity greater than 1.7 on the Pauling Scale.

10. Lattice Structure The ions within an ionic compound arrange themselves within a crystal structure based upon electrostatic forces. ◦ Like charges repel. ◦ Opposite charges attract. The lattice contains a large number of ions. The chemical formula for an ionic compound is really a ratio rather than a true expression of the number of ions within the compound.

11. Lattice Structure Ionic Lattice- Predictable three dimensional crystalline structure. Lattice Enthalpy- The strength of the force between ions within the lattice. Coordination Number- The number of ions which surround a given ion within the lattice.

12. Covalent Bonding: IB Objectives 4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and a positively charged nucleus. 4.2.2 Describe how the covalent bond is formed as a result of electron sharing. 4.2.3 Deduce the Lewis structures of molecules and ions for up to four electron pairs on each atom. 4.2.4 State and explain the relationship between the number of bonds, bond length, and bond strength. 4.2.5 Predict whether a compound of two elements would be covalent from position of the elements on the periodic table or from their electronegativity values. 4.2.6 Predict the relative polarity of bonds from electronegativity values 4.2.7 Predict the shape and bond angles for species with four, three, and two negative charge centers on the central atom using the valence shell electron pair repulsion theory (VSEPR). 4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities. 4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite, and C60 fullerene) 4.2.10 Describe and compare the structure of and bonding of silicon and silicon dioxide

13. Covalent Bonding Covalent bonding involves the sharing of electrons between atoms. Shared electrons are electrically attracted to both nuclei and hold them together via this electrostatic force. Molecule- compound containing covalent bonds. Octet Rule- Tendency of atoms to form stable arrangement of eight outer shell electrons. ◦ Bonding Pairs ◦ Non-Bonding/Lone Pairs

14. Multiple Bonds More that one pair of electrons may be shared to obtain an octet. Double Bond ◦ Two pairs of electrons are shared ◦ Contains one sigma bond and one pi bond Triple Bond ◦ Three pairs of electrons are shared ◦ Contains one sigma bond and two pi bonds

15. Lewis Structures Lewis Structures are a simple notation used to represent the valence electrons of the atoms within a molecule.

16. Lewis Structures: Ions

17. Exceptions to Octet Rule Some elements are stable with less than eight valence electrons (incomplete octet). ◦ Typically involves small central atoms (Be, B). Some elements are stable with more than eight valence electrons (expanded octet). ◦ Involves large central atoms (3 rd period +)

18. Dative Bonds Both of the shared electrons originated from one of the bonded atoms. Also known as coordinate bonds.

19. Bond Strength Single bond<double bond<triple bond Bond length and strength are inversely related.

20. Polar Bonds A bond will be polar if the difference in electronegativity between atoms involved is between 0.4 and 1.8. Electrons are shared unequally between nuclei.

21. VSEPR Theory Valence Shell Electron Pair Repulsion Theory- electron pairs found in outer energy level or valence shell of atoms repel each other and thus position themselves as far apart as possible. Key Points: ◦ Repulsion applies to bonded and non-bonded pairs. ◦ Double and triple bonds behave as negative charge centers. ◦ Number of negative charge centers determine geometric arrangement of electrons. ◦ Shape is determined by angles between bonded atoms. ◦ Repulsive Forces from strongest to weakest:  Lone-pair to lone-pair  Lone pair to bonding pair  Bonding pair to bonding pair

22. Linear Molecules Two negative charge centers. General Formula: AX2

23. Planar Triangular Molecules Three negative charge centers. General Formula: AX3

24. Tetrahedral Molecules Four negative charge centers. General Formula: AX4

25. Other Molecular Shapes

26. Molecular Polarity Molecules with polar bonds may or may not be polar molecules depending on molecular geometry.

27. Crystalline Solids Allotropes of Carbon. Contain only covalent bonds. Have different characteristics. See p. 124.

28. Intermolecular Forces: IB Objectives 4.3.1 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles, or hydrogen bonding) and explain how they arise from the structural features of the molecules. 4.3.2 Describe and explain how intermolecular forces affect the boiling points of substances.

29. Intermolecular Forces Intermolecular Forces exist between the molecules of a substance. These forces depend upon the structure of the molecules involved.

30. Van der Waals’ forces

31. Dipole-Dipole Interactions Attractive forces occurring between partial positive and negative ends of molecules with permanent dipoles.

32. Hydrogen Bonding Special type of dipole-dipole force which occurs between molecules in which hydrogen is bound to extremely electronegative element (F, O, N)

33. Metallic Bonding: IB Objectives 4.4.1 Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons. 4.4.2 Explain the electrical conductivity and malleability of metals.

34. Metallic Bonding Electrons are delocalized from origin metal atom. Metal atoms become cations surrounded by delocalized electrons. Strong attractive force between metal cations and delocalized electrons. Accounts for all characteristics of metals.

35. Physical Properties: IB Objectives 4.5.1 Compare and explain the properties of substances resulting from different types of bonding.

36. Physical Properties Properties which may be observed without chemically altering substance. ◦ Melting/Boiling Points ◦ Solubility ◦ Electrical Conductivity

37. Melting and Boiling Points Melting and boiling points increase with: ◦ Increasing molecular weight ◦ Increasing intermolecular interactions (due to extent of polarity of molecule)

38. Solubility “like dissolves like” Ionic compounds and polar molecules dissolve in polar solvents but not in nonpolar solvents. Non polar molecules dissolve in nonpolar solvents, but not in polar solvents.

39. Electrical Conductivity Ionic compounds can only conduct electricity in aqueous solutions. Metals can conduct electricity due to delocalized electrons. Most covalent molecules do not conduct electricity. Some polar covalent molecules may dissolve, ionize, and conduct.