Chapter 7 “Chemical Formulas and Bonding”

1. Chapter 7“Chemical Formulas and Bonding” “How it all sticks together….” T. Witherup 11/06

2. Some Questions to Consider…. ??? ??? • Why are so few elements (such as Au, Ag, S, N, O) found in Nature in their free atomic state? • Why do atoms of different elements react to form compounds? • What is happening in this process? • How can we explain the millions of compounds that are known today? • Answers to these questions will be found in Chapter 7 (“Chemical Formulas and Bonding”).

3. Chapter 7 Objectives • Describe the characteristics of an ionicbond. • State and use the “Octet Rule.” • Learn how to use “Lewis Dot” diagrams. • Learn the types of ions. • Describe the characteristics of a covalent bond. • Describe the difference between ‘polar’ and ‘non-polar’ covalent bonds. • Write names for ionic compounds, molecular compounds and acids.

4. 7-1 Ionic Bonding • What’s an ‘ION’? • An atom or group of atoms having a charge. • Do you remember how ions form? • Metalslose electrons to become positive ions, called cations. (Which electrons do they lose?) • M  M1+ + e1- • Nonmetalsgain electrons to become negative ions, called anions. (Where do the new electrons go?) • X + e1-  X1-

5. 7-1 Ionic Bonding (cont’d) • Positively charged ions are attracted to negatively charged ions. • Why? • Because ‘opposites attract.’ • Ionic Compound: A substance that is composed entirely of ions. • An ionic formula is the simplest whole-number ratio of the ions, so the total charge balances to zero. • Total (+) charges & total (-) charges = Zero

6. Ionic Compound (General Example) + - and combine to form an Ionic Compound Cations Anions These ions are held together in a solid by electrostatic attraction: - + - + - + - + + - + - + - + - - + - + - + - +

7. Ionic Bonding (Specific Example) • Sodium (Na) is a poisonous, very reactive metal. • Chlorine (Cl2) is a poisonous, very reactive nonmetal. • They combine violently to form ordinary table salt, NaCl, which is relatively harmless. • NaCl is composed of Na1+ and Cl1- ions. • Na  Na1+ + e1- • Cl + e1-  Cl- • Overall: Na + Cl  NaCl (Note the 1:1 ratio.)

8. The Octet Rule • Atoms tend to gain, lose or share electrons in order to acquire a full set (8) of valence electrons. Na = [Ne]3s1 Loses a 3s1 electron to form Na1+ ([Ne] electron core). 1+ (Na1+) e1- e1- 1- Gains an electron in 3p to form Cl1- (3s23p6) ([Ar] electron core). Cl = [Ne]3s23p5 e1- e1- e1- e1- (Cl1-) e1- e1- e1- e1- e1- e1- e1- e1- e1- e1-

9. The Role of Valence Electrons • Note that only the valence electrons were involved in this change, NOT the core electrons. • Why? (Which orbitals & electrons are encountered first when two atoms interact?) • Chemists focus on the valence electrons (outer electrons) to understand the chemistry of atoms. • To aid us, we use shorthand diagrams, called Lewis Dot Diagrams, where dots represent the valence electrons around an atom. • Let’s do some examples.

10. Lewis Dot Diagram Method • Write the element symbol. • Use dots to show the valence electrons (alone or in pairs) around the symbol. • Sodium would be Na with one dot. • Chlorine would be Cl with seven dots. • Our previous reaction of sodium with chlorine would be written as Na. + .Cl:  Na..Cl:  Na1+ + .Cl:1- ¨ ¨ ¨ . ¨ ¨ ¨

11. Lewis Dot Diagrams (Practice) Practice doing this! Remember, show only the valence electrons.

12. Types of Ions • MonoatomicCations • Na1+, Mg2+, Al3+ • Fe2+ [Iron(II)], Fe3+ [Iron(III)] • MonoatomicAnions • F1-, Cl1-, Br1- • PolyatomicIons • NH41+, OH1-, NO31-, SO42-, CO32-, PO43- • See list of ions you MUST learn! • Pages 231 & 232 • http://www.ausetute.com.au/wriiform.html

13. Facts About Ionic Compounds • Binary Ionic Compound - contains ions of only two elements. (e.g. NaCl, CaBr2) • Empirical Formula – the formula of a compound with the lowest whole-number ratio of the elements. • NaCl (NOT Na2Cl2 orNa3Cl3 orNa100Cl100) • The “net charge” of a neutral compound must equal zero, which tells us the ion ratio. (Ca2+ & Cl1- needs CaCl2 as the correct formula.)

14. Rules for Writing Ionic Formulas • Use the simplest whole number ratio of Cation and Anion. • Since the net charge must be zero, balance the number of cations and anions so the total positive charge equals the total negative charge. • Use subscripts after each ion to indicate how many are present. (Omit ‘1’ though.) • Use parentheses around polyatomic ions and indicate their number with a subscript outside the parenthesis. • Crisscross method helps write ionic formulas. • See the next slides.

15. Crisscross Method for Writing Ionic Compound Formulas • Ionic compounds must have a net ionic charge of zero (neutral). • The total + and – charges must cancel. • Always keep polyatomic ions intact! • Use ‘crisscross’ method to write formulas. • The charge superscript becomes the subscript of the opposite ion, indicating the number of ions. • Ba2+ & Br1- becomes BaBr2 [2+ with 2(1-)] = 0 • Al3+ & NO31- becomes Al(NO3)3 [3+ with 3(1-)] = 0 • NH41+ and SO42- becomes (NH4)2SO4 [2(1+) with 2-] = 0

16. Crisscross Method Examples Barium bromide: BaBr2 Ba2+ Br1- becomes Aluminum nitrate: • Al3+ NO31- Al(NO3)3 becomes Notice that ‘1’ is not written, that the nitrate ion is kept intact, and that the net charge is zero. (For example, barium bromide, 1(2+) + 2(1-) = 0)

17. Crisscross Method More Examples Ammonium sulfate: (NH4)2SO4 NH41+ SO42- becomes Notice how the parentheses are used. Aluminum oxide: O2- Al2O3 Al3+ becomes Notice how the net charge is zero. [2(3+) + 3(2-) = 0 PRACTICE, PRACTICE, PRACTICE!

18. Naming Ionic Compounds • Chemists name compounds on the basis of the atoms and bonds present. • Ionic compounds are named from their elements or polyatomic ions. • Cations (+) are named first (usually an element name). • If it can have more than one charge, use Roman numerals to indicate which ion is actually present. • FeCl3 is iron(III) chloride & FeCl2 is iron(II) chloride. • Change the ending of the anion to ‘ide’ (unless a polyatomic ion is present). • NaCl is sodium chloride. • Al2O3 is aluminum oxide. • Ba(NO3)2 is barium nitrate. • K2SO4 is potassium sulfate. • What is NiBr2? Sr3(PO4)2? FeI2?

19. Hydrates • Hydrate – Ionic compound that absorbs water into their crystals. • Blue copper sulfate contains several water molecules in its crystal. We will do a lab about this. • Anhydrous – A water-free substance. • These ionic compounds are named to reflect the ‘water of hydration.’ • Name the compound in the normal way. • Add the word ‘hydrate’ and a prefix term to show the number of water molecules (degree of hydration). • See Fig. 7-24 on page 246. • Di-, tri- tetra-, penta- etc. • MgSO4 *7 H2O is magnesium sulfate heptahydrate. • What is the formula for copper(II) sulfate pentahydrate?

20. Properties of Ionic Compounds • High melting points (usually). • NaF (996 °C), NaCl (801 °C) • This indicates very strong ionic bonding. • Very brittle. • Shatter, or cleave, in fixed paths rather than randomly. • Example: Rock salt. • Water soluble (usually). • Water breaks the ionic bonds. • Aqueous solutions conduct electricity because the ions are free to move about in the water. • Conduct electricity when molten (liquid). • Ions are freed from the crystal structure (lattice). • Do not conduct electricity when solid. • Ions are held firmly in place, so they simply vibrate.

21. 7-2 Covalent Bonding • A covalentbond is formed by a shared pair of electrons between two atoms. • Molecule – group of atoms united by a covalent bond. • MolecularSubstance – a material made up of molecules. • Empirical Formula - the formula of a compound with the lowest whole-number ratio of the elements. • Molecular Formula – chemical description of a molecular compound or molecule. • Structural Formula – a formula that specifies which atoms are bonded to each other in a molecule. • Lewis Structures – molecular structure based on Lewis Dot diagrams.

22. Covalent Bond Formation Sharing of electrons, as in two chlorine atoms! .. .. combines with .Cl: :Cl. .. .. to form a Cl2 molecule by sharing electrons. .. .. :Cl:Cl: .. .. This is a “diatomic” molecule, along with molecules of fluorine, bromine, iodine, hydrogen, nitrogen, and oxygen. “Professor BrINClHOF” will help you remember them!

23. Describing Covalent Bonds • Draw Lewis dot diagrams, including unshared pairs of electrons. • Use a ‘dash’ for each pair of electrons in a bond. • Examples: Chlorine (Cl2) is written as Cl-Cl. • Single covalent bonds • C:C or simply C-C (Note the ‘dash.’) • Double covalent bonds • C::C or simply C=C (Note the ‘double dash.’) • Triple covalent bonds • C:::C or simply CΞC (Note the ‘triple dash.’)

24. Properties of Covalent Compounds • Low melting points (usually). • Methane, (CH4) is a gas at room temperature; oils are liquids at room temperature; wax melts at ~100°C. • This indicates very weak molecular association. • Soft. • Wax feels slippery and may be deformed even as a solid. • Insoluble in water (usually). • Water cannot break the covalent bonds. • Aqueous solutions do not conduct electricity (no ions are free to move about in the water). • Do not conduct electricity when molten (liquid). • Again, there are no ions to move about. • Do not conduct electricity when solid. • No ions!

25. Properties of Covalent Bonds (cont’d) • Remember “electronegativity”? (What is it?) • The ability of an atom to attract electrons in a chemical bond. • Fr has the lowest (0.7) and F has the highest (4.0) on the Pauling scale. • Electronegativity differences (“delta EN” or ∆EN) dictate which atom in a bond more strongly attracts the electrons. • See Fig 7-20, page 242, and the following slide. • Chemists use lower case Greek letter delta (δ) to mean a “partial” or “small difference.”

26. Polarity • Refers to the unequal sharing of electrons in covalent bonds of compounds. • When both atoms in a bond are identical, they form NONPOLAR bonds (e.g. Cl2 or F2) because there is, equal sharing. • When one atom has higher electronegativity than the other, it forms a POLAR bond (e.g. HCl), which means the electrons are not shared equally. • We use delta +/- (δ+ or δ-) or arrows (+) to show polarity of a bond. H-Cl |

27. Bond Type by Electronegativity(Use the electronegativity difference, ∆EN, to predict the bond type.) Note that a large∆EN means that it is an ionic bond. Electrons have transferred from one atom to another.

28. A Special Type of Bonding • Metallic Bonding – the force of attraction that holds metals together. • Positive metal ions are in a ‘sea of electrons’ (freely floating valence electrons) that are shared. • This accounts for metallic properties, such as electrical conductivity, luster, ductility, malleability. • Drifting electrons insulate the metal ions from one another, so the ions can easily slide past each other when stressed, unlike ionic solids, which shatter when stressed.

29. Exceptions to the Octet Rule • Atoms with less than an octet. • Boron compounds. • Atoms with more than an octet. • Atoms with d-electrons, such as sulfur. • Molecules with an odd numberof electrons. • So called “Radicals” like nitroxyl, NO.

30. 7-3 Naming Chemical Compounds • Ionic compounds are named from their elements or polyatomic ions. • Hydrates have water in their solid structure, but anhydrous substances do not. • Molecular compounds are named using prefixes to indicate the number of atom in the formula. • Acids have special names that must be memorized (Fig 7-27, pg 249). • PRACTICE, PRACTICE, PRACTICE!

31. Naming Molecular Compounds • Use the element names and prefixes to indicate the number of atoms in the formula. • Di-, tri-, tetra-, etc. • CO is carbon monoxide. (“Mono” is not used for the first element generally.) • CO2 is carbon dioxide. • N2O is dinitrogen monoxide. • N2O4 is dinitrogen tetroxide. (Not usually ‘tetraoxide’ because it is hard to say!) • Name these: N2O5. SO3. BF3. PF5 • Many molecular compounds have ‘common names.’ • Dihydrogen monoxide is ______? • Trihydrogen mononitride is ‘ammonia.’

32. Naming Common Acids • Acids are molecular substances that dissolve in water to produce hydrogen ions (H+). • Acids have special names that must be memorized (Fig. 7-27, page 249), but focus on these and their anions: • Hydrofluoric, hydrochloric, hydrobromic, hydroiodic, • Nitric • Sulfuric • Carbonic • Phosphoric • Acetic

33. Did we meet the Chapter 7 Objectives? • Describe the characteristics of an ionic bond. • State and use the “Octet Rule.” • Learn how to use “Lewis Dot” diagrams. • Learn the types of ions. • Describe the characteristics of a covalent bond. • Describe the difference between ‘polar’ and ‘non-polar’ covalent bonds. • Write names for ionic compounds, molecular compounds and acids.