Valence Electrons and Chemical Bonding

1. Unit 3 IONIC BONDS COVALENT BONDS Slides 18 (MOLECULAR COMPOUNDS) – 47 (MOLECULAR PROPERTIES) Sharing is caring. CHEMICAL FORMULAS • Slides 2 (IONS) -17(ALLOYS) • Bond. Ionic bond. Taken, not shared. • Slides 48 (WHAT’S IN A NAME?) – 58 (NAMES & FORMULAS FOR BASES)

2. Ions • Valence electrons are the electrons in the highest occupied energy level of an element’s atoms. • They are largely responsible for the chemical properties of an element. • They are also the primary reason for similarities in groups/families of elements.

3. Valence Electrons • Group numbers, i.e. 1A(1), 2A(2), 3A(13), etc. correspond to the number of valence electrons of that group. 1A = 1 valence electron, 2A = 2 valence electrons, etc. For groups 13-18, subtract 10. • Exception – Helium has only 2 valence electrons even though it is in group 8A. • Valence electrons are usually the only electrons used in chemical bonds.

4. Electron Dot Structure Electron dot structures are diagrams that show valence electrons as dots.

5. Octet Rule • Atoms of metals tend to lose their valence electrons, leaving a complete octet in the next-lowest energy level. • Atoms of some nonmetals tend to gain electrons or share electrons with another nonmetal to achieve a complete octet.

6. Forming Cations • An atom’s loss of valence electrons produces a cation, or a positively charged ion. • The most common cations are produced by the loss of valence electrons from metal atoms. • Some ions formed by transition metals do not have noble-gas configurations, and are exceptions to the octet rule. • Cations have the same name as the atoms from which they form. For example, Na is neutral and Na+, is simply the sodium cation.

7. Forming Anions • The gain of electrons by a neutral atom produces an anion. • Unlike cations, anions do not have the same name as their neutral counterparts. When a nonmetallic element forms an anion, the suffix –ide is added to the name. • For example Cl- is the chloride ion.

8. Anions • The ions that are produced when atoms of halogens gain electrons are called halide ions. • Example • F- = fluoride ion • Cl- = chloride ion • Br- = bromide ion

9. Ionic Bonds and Ionic Compounds • Compounds composed of cations and anions are called ionic compounds. • An example would be NaCl, sodium chloride, or table salt. Na+ combines with Cl-. • Although they are composed of ions, ionic compounds are electrically neutral. The positive cations equal the negative anions.

10. Ionic Bonds • Because of their opposite charges, cations and anions attract each other using electrostatic forces. • The electrostatic forces are called ionic bonds and are what holds the ions together in an ionic compound.

11. Formula Units • Chemical formulas are used to represent the kinds and numbers of atoms in the smallest representative unit of a substance. • For example, NaCl is the smallest representative unit of the sodium chloride compound. Ionic compounds exist as a collection of repeating units. The chemical formula is simply a ratio known as a formula unit.

12. Formula Units • A formula unit is the lowest whole-number ratio of ions in an ionic compound. • NaCl is a 1:1 ratio. • Mg2+ and Cl- have a ratio of 1:2 because its formula unit is MgCl2.

13. Properties of Ionic Compounds • Most ionic compounds are crystalline solids at room temperature. • Component ions in such crystals are arranged in repeating 3-D patterns. • Ionic compounds generally have high melting points.

14. Properties of Ionic Compounds • Coordination numbers of ions are the number of ions of opposite charge that surround the ion in a crystal. • For example, in a crystal of NaCl, the sodium cation is surrounded by six chloride ions; thus Na+ has a coordination number of 6. • Ionic compounds can conduct an electric current when melted or dissolved in water.

15. Bonding in Metals • Metals are made up of closely packed cations rather than neutral atoms. • The valence electrons of metal atoms can be modeled as a sea of electrons. • Metallic bonds consist of the attraction of the free-floating valence electrons for the positively charged metal ions, resulting in the bonds that hold metals together.

16. Alloys • Alloys are mixtures composed of two or more elements, at least one being a metal. • Alloys are important because their properties are often superior to those of their component elements.

17. Alloys • Alloys can form from their component atoms in different ways. • Component atoms of relatively the same size, can replace each other in the crystal; this type of alloy is called a substitutional alloy. • If atomic sizes differ greatly, smaller atoms can fit into interstices (spaces) between larger atoms; this alloy is called an interstitial alloy. • Steels are examples of interstitial alloys.

18. Draw the Lewis Dot Structures for: • Nitrogen • Oxygen • Fluorine • Carbon • Silicon • Phosphorus • Sulfur • Chlorine • Hydrogen

19. Target: Understanding the concepts that are a part of covalent bonding • Sharing compared to Giving / Taking • Lone pair / shared pair / lone electron • Lewis dot – octet rule • Single / double/ triple bonds • Monatomic / Diatomic • Which atoms • Salt / molecule • Melting / boiling points • Molecular / Structural Formula

20. Molecular Compounds • Ionic compounds deal with the loss or gain of electrons in their bonds. • Covalent bonds are held together by the sharing of electrons.

21. Atoms • Some elements are monatomic, a single atom in nature. • Examples are helium and neon. • Some elements are molecules in nature. • A molecule is a neutral group of atoms joined together by covalent bonds. • A diatomic element is a molecule consisting of two atoms. • Hydrogen, oxygen, nitrogen and all of the halogens are diatomic in nature. (The “Gens”) • Diatomic molecules do not have to be composed of the same element.

22. Molecular Compounds • A compound composed of molecules is called a molecular compound. • The molecules of a particular molecular compound are all the same. • Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds. The exception would be network solids. • Many are gases or liquids at STP. • A molecular formula shows how many atoms of each element a molecule contains.

23. Molecular Formulas • A formula tells you how much and what type of element(s) is/are in a molecule it does not tell you about a molecule’s structure. • Different types of models can be used to physically illustrate the molecular formula and structure of molecular compounds. CH4 (methane) H2O (water)

24. The Nature of Covalent Bonding • In covalent bonds, electron sharing usually occurs so that atoms attain the electron configuration of noble gases. • Combinations of atoms of nonmetals and metalloids in Groups 14(4A) – 17(7A) are likely to form covalent bonds. • In this way they form an octet and achieve noble gas configuration.

25. Single Covalent Bonds • Two atoms held together by the sharing of a pair of electrons are joined by a single covalent bond. • An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots.

26. Structural Formula • A structural formula represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms. • Example: H – H, where the dash represents the two shared electrons.

27. Covalent Bonding • A pair of valence electrons that is not shared between atoms is called an unshared pair, a lone pair or nonbonding pair. • Example: Notice that only 2 of fluorine's electrons are shared in the covalent bond. The other pairs surrounding the fluorine atoms are unshared pairs. In general only unshared electrons will be used for bonding.

28. Quiz today over yesterday’s notes: • Get Your Clickers • You have 5 minutes to review before we start. • No talking during quiz

29. What is the difference between ionic and covalent bonding • Ionic shares and covalent gives / takes • Ionic has metal/nonmetal ionic bonds, covalent has nonmetal/nonmetal metallic bonds • Ionic gives/takes covalent shares • The only difference between ionic and covalent are the melting points

30. Which electrons bond in covalent bond • Lone pairs • Shared pairs • Unshared pairs • Lone electrons

31. If there are 4 lone electrons, how many bonds will that atom make • 1 • 2 • 3 • 4

32. What is a covalent compound called? • Salt • Molecule • Empirical • Molar

33. If only two atoms make a covalent bond and they both have only 6 electrons, what type of bond will they make? • 2 doubles • 1 Double • Triple and a single • 4 singles

34. Target: Understanding the concepts that are a part of covalent bonding • Sharing compared to Giving / Taking • Lone pair / shared pair / lone electron • Lewis dot – octet rule • Single / double/ triple bonds • Monatomic / Diatomic • Which atoms • Salt / molecule • Melting / boiling points • Molecular / Structural Formula

35. Target: Understanding the concepts that are a part of covalent bonding • Double & Triple bonds • Coordinate Covalent Bonds • Review Polyatomic Ions • Resonance • Octet Rule Exceptions • Molecular Orbitals • Sigma & pi bonds

36. Double & Triple Covalent Bonds • Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons. • A bond that involves two shared pairs of electrons is a double covalent bond. • A bond formed by sharing three pairs of electrons is a triple covalent bond.

37. Double Bonds • Carbon dioxide (CO2) forms 2 double bonds; one between carbon and each oxygen. O=C=O • CO2 is commonly found naturally in the atmosphere and is also used to carbonate soft drinks and other beverages. • In double bond (C=O), Oxygen is stable, but carbon is not. If oxygen donates a PAIR of unshared electrons then both are stable CO.

38. Coordinate Covalent Bonds • Carbon monoxide (CO) exists in a different type of covalent bond. • Carbon needs to gain 4 electrons to achieve neon’s noble gas configuration and oxygen needs 2 electrons. • Both achieve noble gas configurations by a type of bonding called coordinate covalent bonding. • A coordinate covalentis formed when one atom contributes both bonding electrons in a covalent bond. Example: Carbon monoxide

39. Polyatomic Ions • An ammonium ion (NH4+) consists of atoms joined by covalent bonds, including a coordinate covalent bond. • A polyatomic ion such as NH4+ is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit.

40. Ammonia to Ammonium • This ion forms when a positively charged hydrogen ion (H+) attaches to the unshared electron pair of an ammonia molecule (NH3)

41. Polyatomic Ions • Most polyatomic cations and anions contain both covalent and coordinate covalent bonds. Therefore compounds containing polyatomic ions include both ionic and covalent bonding. • The electron dot structure for a neutral molecule contains the same number of electrons as the total number of valence electrons in the combining atoms. • Negative charges on a polyatomic ion shows the number of electrons in addition to the valence electrons of the atoms present.

42. Bond Dissociation Energies • The energy required to break the bond between two covalently bonded atoms is known as the bond dissociation energy. • A large bond dissociation energy corresponds to a strong covalent bond.

43. Quiz on Covalent / Ionic Bonding • Determine the type of bond from the elements given. Determine the balanced formula. Draw Lewis Dot structure for the compound. • P and F • N (bonding with itself) • Mg and Cl

44. Agenda • Naming molecular bonds (suffixes and all)

45. What’s in a name? • Ions of Transition Metals • Many of the transition metals 3-12 (1B-8B) • The charges of the cations of many transition metal ions must be determined from the number of electrons lost.

46. Naming Ions • There are 2 methods for naming transition metal cations. • The preferred method is called the Stock system, in which a Roman numeral representing the charge of the transition metal is placed in parentheses. • Example: Fe(II)S • The older, less useful method uses a root word with different suffixes at the end of the word. • Example: ferrous sulfide • There are few transition metals that have only one charge; Ag+, Cd2+, Zn2+. (You must memorize these.)

47. Writing Formulas for Ionic Compounds • Binary compounds are compounds composed of two elements and can be either ionic or molecular (covalent). • If you know a formula you can write its name. • First verify that the compound is composed of a monatomic metallic cation and a monatomic nonmetallic anion. • To name a binary ionic compound, place the cation name first, followed by the anion name. • Remember, with anions, the –ides have it. • Write the symbol of the cation and then the anion. Add whatever subscripts are needed to balance the charges. • Remember to follow the “zero” rule. • Figure out the charge you know, follow the zero rule to find the charge you don’t know. • Following the crisscross method will help you balance, 9 times out of 10. • Example: Fe3+ and O2- • Combine by taking 2 from oxygen and using it as iron’s subscript. Then take 3 from iron and use as oxygen’s subscript. Fe2O3

48. Polyatomic Ions • Composed of more than one atom. • The names of most polyatomic anions end in –ite or –ate. • Example: (ClO-) is the hypochlorite ion (CO32-) is the carbonate ion. • Most are negative polyatomic anions, while a few are positive polyatomic cations.

49. Compounds with Polyatomic Ions • Remember that if it ends with –ite or –ate, it contains a polyatomic anion that includes oxygen. • Write the symbol for the cation followed by the formula for the polyatomic anion and balance the charges. • Be sure the subscripts are in the lowest whole number ratio. • Parentheses must be placed around polyatomic ions that have a subscript for balance. • Example: Ca2+ combined with NO3- would be written as Ca(NO3)2 • The same rules apply to polyatomic cations, such as NH4+. • Write the polyatomic cation first, then follow with the anion. • Use parentheses if a subscript is added to the polyatomic cation for balancing. • Example: NH4+ combined with S2- would be written as (NH4)2S

50. Naming & Writing Formulas for Molecular Compounds • A prefix in the name of a binary molecular compound tells you how many atoms of an element are present in each molecule of the compound. • Example: mono – one, as in monorail, di – two, as in dissect, tri – three, as in tricycle. • Note: prefixes are NEVER used in ionic compounds