Ionic and Covalent bonds

1. Ionic and Covalent bonds Joe Glavan Riverside Local School District December 17, 2013

2. Chemical bonding • Chemical bond- is a mutual electrical attraction between the nuclei and valance electrons of different atoms that binds the atom together. • Two types of bonds; ionic and covalent

3. Two types • Ionic • Cations to anion (metal to nonmetal) • Covalent • Sharing of electrons (nonmetal to nonmetal)

5. Bonding is never purely ionic or covalent • Bonding is never purely ionic or covalent • It depends on the ability to attract electrons. • Which is ___________________.

6. Electronegativity • We can use the electronegativity difference of the atom to determine the bond type. • Using the electronegativity table: • 3.3 – 1.7 ionic • 1.7 – 0.3 polar covalent • 0.3 – 0.0 Nonpolar covalent • What is the electronegativity trend?

8. Section 1: Bonding and Covalent Bonding

9. Section 1: Covalent Bonding • The difference in electronegativities determines where the bonding electrons spend their time, and therefore the bond type Eneg diff = 1.7 – 3.3 Eneg diff = 0 – 0.3 Eneg diff = 0.3 -1.7

11. Covalent Bonds • What is polar and nonpolar covalent mean? • It is the placing of the shared electron, if one atom pulls the electron closer to it, then you get polar covalent. If they are held close to both then it is nonpolar. • Polar – meaning one element is pulling the electron more then the other. Making a positive and a negative end • Nonpolar – equal sharing (no positive and negatives ends) usually the same element bond to itself.

15. Ionic compounds • Metal plus nonmetal • Cation plus anion • + ion plus – ion • Remember that the positive charge must cancel the negative charge

16. Formula Units • Simplest collection of atom from which an ionic compound’s formula can be established. • Remember that the overall charge on a ionic compound is neutral.

17. Examples • Na+ -- Cl- ----- NaCl • Notice that the charges are cancelled out. • The name is sodium chloride • Ca+2 F----- CaF2 • Notice that the charges cancel out because 2 F and 1 Ca. • Ca +2 F - F -

18. Ionic compounds • To write a formula unit you must follow these rules: • 1. The symbol of metal or cation is written first • 2. The symbol of the anion is written second • 3. The charges are either criss-crossed and make into subscripts or more atoms are added until the overall charge is zero. Ca+2 Cl- Cl- +2 -2 = zero

19. Formula Writing with Ionic Compounds • Recall that polyatomic ions are groups of nonmetals with an overall charge. These will be criss-crossed and canceled out as a group. • Parentheses are needed when you need multiple polyatomic ions. • Mg+2 PO4-3 • Mg3(PO4)2 • You need two phosphates for three magnesium in order to have an overall charge of zero.

20. Practice • Li and F • Na and Br • Ca and I

21. Lewis structures • Draw the lewis structure for Li • Draw the lewis structure for F • Show the transfer of electrons

22. Ionic Characteristics • Ionic compounds will often form crystal lattice structures. Purple is Sodium Na Green is Cl

23. Ionic Characteristics • Bond Strength in ionic compounds, amount of energy released when separated ions in a gas come together to form a crystalline solid. • Lattice energy is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. • Strong bonds, Increased melting/boiling point (Do not vaporize as readily as covalent)

24. Ionic bonds are strong and have the following properties • High melting and boiling points • Hard • Brittle • In a molten state, or when dissolved in water ionic compounds are electrical conductors, because the ions can move freely to carry the electrical charge. • (solid state- no movement- no electrical conductivity) • Dissolved in water – electrical conductors.

25. Why are ionic compounds hard but brittle? • In an ionic crystal, even a slight sift of one row of ions relative to another causes a large build up of repulsive forces. These forces make it difficult for one layer to move relative to another, causing ionic compounds to be hard. If one layer is moved, however the repulsive forces make the layer part completely, causing ionic compounds to be brittle.

26. Strong but brittle

27. Polyatomic Ions • Atoms that are covalently bonded to each other to form a group of atoms that has both molecular and ionic characteristics. • The charge on the polyatomic ion results from an excess of electrons, or shortage. • Polyatomic (-) plus a metal • Polyatomic (NH4+ ) plus a nonmetal

28. Practice • Li + F • S + Ca • Ba + N • Na + CO3-2 • NH+ + Cl

29. Naming Binary Ionic Compounds from Empirical Formulas • Ionic compounds - named by stating the metal or ammonium first • Take the root of the second non-metal element and add the suffix –ide.

30. Name Me • NaCl • CsI • CaBr2 • Potassium Oxide • Potassium Phosphide

31. Naming Polyatomic Ionic Compounds • Polyatomic ionic compounds - named by stating the metal or the + polyatomic ion first • Then state the name of the negative polyatomic ion.

32. Name Me • Al2(SO3)3 • (NH4)2CrO4 • Sodium Carbonate • Ammonium Chloride

33. Writing Names of Ionic Compounds with Multivalent Metals • Write the names of the ions in the order they are written. • Reverse criss cross to determine the charge of the multivalent metal • Remember that some ratios have been reduced • Write the charge in Roman numerals in parentheses after the metal

34. Name Me • Iron (II) Oxide • CoO • MnO2 • Au2(CrO4)3 • Iron (II) Phosphate

35. Give the Formula For • Iron (II) Chloride • Titanium (IV) Oxide • Chromium (III) Acetate • Lead (II) Oxide • Gold (III) Oxalate

36. METALLIC BONDING • How does metallic bonding compare to covalent and ionic. • Characteristics of metallic bonds • Predict properties based on the bond type

37. Metallic bonds • Chemical bonding that results from the attraction between metal atoms and surrounding sea of electrons • Sea of electrons is due to the fact that there are overlapping empty orbital and the delocalized electrons can move freely throughout the entire metal. • Electrically charged

38. Metallic Bonding

39. What are the properties of metals? • Why do they have these properties

40. Metallic Properties • High electrical conductor • High thermal conductor (free e-) • Absorb a wide range of light frequencies because of the small energy differences between orbital's. – makes them shinny • Malleability – ability of a substance to be hammered into a thin sheet • Ductility – ability of a substance to be drawn into a wire. • This is possible because the metallic bonding is th same in all directions throughout the solid. One plane can slide past the other plane without encountering any; resistance or breaking any bonds.

41. Metallic Bond Strength • Metallic bond strength varies with the nuclear charge of the metal atoms and the number of electrons in the metal’s electron sea. • Reflected in the heat of Vaporization. When a metal is vaporized, the bonded atoms in the normal (usually solid) state are converted to individual metal atoms in the gaseous state.

42. Covalent Bonding • Ionic substances have such characteristics as being brittle, crystalline, and having high melting/boiling points (solid at room temp) • Many other substances have different properties such as water, oxygen and carbon dioxide, which are all liquids or gases at room temperature – each of these have low melting/boiling points • These properties result from a different type of bonding in which no electrons are transferred, but rather a pair or pairs of electrons are shared by atoms to form covalent bonds

43. Covalent Bonding • In covalent bonding, two atoms approach each other and begin to experience both attractive and repulsive forces • The atoms settle in at a characteristic distance (bond length) at which their potential energies are at a minimum

44. Covalent Bonding • The potential energy possessed by the arrangement of atoms when they are at the bond length is the bond energy • Bond energy is the amount of energy needed to break the bond, and is used to measure the strength of the bond

45. Covalent Bonding • In covalent bonds, each atom still achieves an octet of electrons, but now it is accomplished by sharing rather than by electron transfer • For example, when two fluorine atoms bond to form F2

46. Covalent Bonding • The bond energies and lengths of single bonds, double bonds, and triple bonds are, shockingly, not the same

47. Covalent Bonding • The bond energies and lengths of single bonds, double bonds, and triple bonds are, shockingly, not the same

48. Covalent Bonding • Lewis structures are a representation of a covalent compound • To draw a Lewis structure:1. Determine the total number of valence e- from all atoms2. Determine the central atom * C is always the central atom, if present * Otherwise, the less electronegative atom is the central atom * H is ALWAYS on the outside of the Lewis structure3. The other atoms present bond to the central atom using a shared pair of electrons4. Add lone/unshared pairs of electrons to all atoms to satisfy theoctet rule – make sure ALL electrons have been used!

49. Covalent Bonding • Use the periodic table to find the valence electrons in an atom!

50. Covalent Bonding • Go through the steps to draw a Lewis structure for H2O