4.1 Ionic Bonding & Structure

1. Mrs. Page IB Chem. 2015-2016 4.1 Ionic Bonding & Structure

2. Understandings • Positive ions (cations) form by metals losing valence electrons • Negative ions (anions) form by non-metals gaining electrons • The number of electrons lost or gained is determined by the electron configuration of the atom • The ionic bond is due to electrostatic attractions between oppositely charged ions. • Under normal conditions, ionic compounds are usually solids with lattice structures.

3. Application & Skills • Deduction of the formula and name of an ionic compound from its component ions, including polyatomic ions • Explanation of the physical properties of ionic compounds (volatility, electrical conductivity, and solubility) in terms of their structure.

4. Nature of Science • Use theories to explain natural phenomenon – molten ionic compounds conduct electricity but solid ionic compounds do not. The solubility and melting points of ionic compounds can be used to explain observations.

5. Introduction to Bonding • Chemical bond: an interaction between atoms or ions that results in a reduction of the potential energy of the system, thereby becoming more stable • Three types of bonds: ionic, metallic, and covalent • The bond type depends on the atoms’ electronegativity values

6. More Bonding Basics • If the atoms have very different electronegativity values, then ionic bonding occurs • If they both have high electronegativity values, then covalent bonding occurs • If they both have low electronegativity values, then metallic bonding occurs

7. Practice: What Kind of Bond? • Na and Cl • Sr and O • C and O • Ni and Fe • N and O • Li and N • Ti and Cr Ionic Ionic Covalent Metallic Covalent Ionic Metallic

8. Valence Electrons • Valence electrons are the electrons in the outermost energy level, which is the highest occupied energy level • They are the electrons responsible for the chemical properties of atoms • Electron transfers result in nobal gas electron configurations. • Core electrons – are those in the energy levels below.

9. Keeping Track of Electrons • Atoms in the same group have the same outer electronic structure and therefore the same number of valence electrons. • The number of valence electrons is easily determined. It is the group number • Group 2: Be, Mg, Ca, etc. • Each has 2 valence electrons

10. Electron Dot (Lewis Dot)diagrams • A way of showing & keeping track of valence electrons. • Write the symbol - it represents the nucleus and inner (core) electrons • Put one dot for each valence electron (8 maximum) • First two electrons placed as a pair (s orbital) • Then they don’t pair up until they have each have 1 (Hund’s rule) X

11. The Electron Dot Diagram (Lewis Structure) for Nitrogen • Nitrogen has 5 valence electrons to show. • First we write the symbol. N • Place 1st pair • Then add 1 electron at a time to each side.

12. The Octet Rule • The noble gases are unreactive in chemical reactions • In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules • The Octet Rule: in forming compounds, atoms tend to achieve a noble gas structure; 8 in the outer level is stable • Each noble gas (except He, which has 2) has 8 electrons in the outer level

13. Formation of Cations • Metals become oxidized (lose electrons) to attain a noble gas structure. • They make positive ions (cations) • If we look at the electronic structure, it makes sense to lose electrons: • Na: 2, 8, 1 shows 1 valence electron • Na1+: 2, 8 This is a noble gas structure with 8 electrons in the outer level.

14. Electron Dots For Cations • Metals have few valence electrons (usually 3 or less); calcium has only 2 valence electrons Ca

15. Electron Dots For Cations Metals will lose the valence electrons Ca

16. Electron Dots For Cations • Form positive ions Ca2+ This is named the “calcium ion”. No dots are now shown for the cation.

17. Electron Configurations: Anions • Nonmetals are reduced (gain electrons) to attain noble gas electronic structures. • They make negative ions (anions) • S = 2,8,6 = 6 valence electrons • S2- = 2,8,8 = noble gas structure. • Halide ions are ions from chlorine or other halogens that gain electrons

18. Electron Dots For Anions • Nonmetals will have many valence electrons (usually 5 or more) • They will gain electrons to fill outer shell. 3- P (This is called the “phosphide ion”, and should show dots)

19. Stable Electron Configurations • All atoms react to try and achieve a noble gas structure. • Noble gases have 8 valence electrons and so are already stable • This is the octet rule (8 in the outer level is particularly stable). Ar

20. Ionic Bonding • Anions and cations are held together by the electrostatic attraction due to the opposite charges (+ and -) • Simplest ratio of elements in an ionic compound is called the formula unit. • The bond is formed through the transfer of electrons (lose and gain) • Electrons are transferred to achieve noble gas structure.

21. Ionic Bonding Cl Na The metal (sodium) is oxidized (loses its one electron) from the outer level. The nonmetal (chlorine) is reduced (gains one electron) to fill its outer level, and will accept the one electron that sodium is going to lose.

22. Ionic Bonding [ ] Cl - Na+ Notes: • Remember that no dots are now shown for the cation • Brackets must be shown in Lewis structures to show overall charge of ion • This step involving the formation of ions is a critical step to show in the formation of ionic bonds

23. Ionic Bond • Negative charges are attracted to positive charges. • Negative anions are attracted to positive cations. • The result is an ionic bond. • A three-dimensional crystallattice of anions and cations is formed.

24. Ionic Compounds • The ionic substance is held together by strong electrostatic attractions in all three dimensions • No molecules present • An ionic lattice-type structure is formed • This gives them distinct physical properties

25. NaCl CsCl TiO2

26. Properties of Ionic Compounds • Hard, brittle crystalline solids • Relatively high melting and boiling points • Do not conduct electricity when solid, but do when molten or in aqueous solution • Are more soluble in water than other solvents

27. - Page 198 The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

28. Preserve Electroneutrality • When ions combine, electroneutrality must be preserved. • In the formation of magnesium chloride, • 2 Cl- ions must balance a Mg2+ ion: • Mg2+ + 2 Cl- → MgCl2

29. Predicting Ionic Charges Group 1A: Lose 1 electron to form 1+ ions K+ H+ Li+ Na+ Rb+

30. Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Sr2+ Ba2+

31. Predicting Ionic Charges Group 13: B3+ Al3+ Ga3+ Loses 3 electrons to form 3+ions

32. Predicting Ionic Charges Nitride N3- Group 15: P3- Phosphide As3- Arsenide Gains 3 electrons to form 3- ions

33. Predicting Ionic Charges O2- Oxide Group 16: S2- Sulfide Se2- Selenide Gains 2 electrons to form 2- ions

34. Predicting Ionic Charges Group 17: F- Fluoride Br- Bromide Cl- Chloride I- Iodide Gains 1 electron to form 1- ions

35. Predicting Ionic Charges Stable noble gases do notform ions! Full octets. Group 18:

36. Predicting Ionic Charges Many transition elements have more than one possible oxidation state. Note the use of Roman numerals to show charges Iron (II) = Fe2+ Iron (III) = Fe3+

37. Polyatomic Ions (MUST KNOW!) • Because these are ions they also are involved in ionic bonding!

38. HOMEWORK • Read pp. 93-97 • Upcoming Dates: December 6 Final Exam Topics 1-3