Chemical Bonding

1. Chemical Bonding • Chemical Bond-force that holds two atoms together. Atoms either transfer electrons and then form ionic compounds or they share electrons to form covalent compounds. • In both cases, the bond forms because of an increase in stability.

2. Chemical Bonding • By forming bonds, atoms acquire an octet of electrons and the stable electron configuration of a noble gas. Electron-dot configuration

3. Formation of Positive Ions • Atoms can lose one or more valence electrons to obtain noble gas electron configuration and form a positive ion called a Cation. • Sodium has one 3s valence electron which it loses to form a Na+ cation.

4. Transition Elements Transition elements have full s orbitals and as you move from left to right on the periodic table electrons are being added to inner d orbitals. Transition metals generally lose the two s orbital electrons to form 2+ cations but in some cases d orbital electrons can be lost to form cations with charges greater than 2+

5. Pseudo-Noble Gas Configuration • Elements from periods 4-6 lose electrons to form outer orbitals with full s, p and d orbitals. Zinc ion formation Zinc ([Ar]4s23d10) Zinc2+ ([Ar]3d10) + 2e- pseudo-noble gas configuration

6. Formation of Negative Ions • Nonmetals have large attraction for electrons and attain stability by gaining electrons to form negative ion called an Anion. ex. Chlorine (1s22s22p63s23p5) Chloride anion (1s22s22p63s23p6) [Ar] valence configuration • Nomenclature for anion is the addition of –ide to name. i.e. Cl-1 is named chloride

7. The charge on the ion is known as the oxidation number of the atom.

8. Formation of Ionic Bond • Ionic Bond- electrostatic force that holds opposite charged particles together in an ionic compound.

9. Formation of Ionic Compound NaCl • Sodium is in Group 1, so it has one valence electron. Chlorine is in Group 17 and has seven valence electrons. • The one valence electron of sodium is transferred to the chlorine atom, both become stable with noble gas electron configurations. • The chlorine atom now has an extra electron and has a negative charge. • Sodium lost an electron and has a positive charge.

10. Ionic Compounds Nomenclature • Most ionic compounds are called salts. • Metal + Oxygen form ionic compounds called oxides.

11. Binary Ionic Compounds • Ionic compounds with only two different elements are termed Binary. Binary ionic compounds can contain more than one ion of each element, as in CaF2, but they are not composed of three or more different elements.

12. Binary Ionic Compounds To name a binary ionic compound, first write the name of the positively charged ion, usually a metal, and then add the name of the nonmetal or negatively charged ion, whose name has been modified to end in -ide. Potassium combines with chlorine to form potassium chloride salt. Magnesium combines with oxygen to form magnesium oxide.

13. Formation of Aluminum Oxide • Aluminum is a Group 13 metal, so it loses its three outer electrons to become an Al3+ ion; oxygen is in Group 16 and has six valence electrons, so it gains two electrons to become an O2- ion. • All the electrons must be accounted for, therefore more than one oxygen atom must be involved in the reaction.

14. Formation of Aluminum Oxide • In all, two Al3+ ions must combine with three O2- ions to form Al2O3. • Remember that the charges in the formula for aluminum oxide must add up to zero.

15. Properties of Ionic Compounds • Chemical bonds between atoms in a compound determine many physical properties of the compound. • In an ionic compound the positive and negative ions are packed into a regular repeating pattern that balances the forces of attraction and repulsion between ions forming Ionic Crystals.

16. Ionic Crystal

17. Properties of Ionic Compounds • Ionic compounds are composed of well-organized, tightly bound ions forming a strong, three-dimensional crystal structure. • Ionic compounds are crystalline solids at room temperature with relatively high melting and boiling points. • In the solid-state ionic compounds are nonconductive due to fixed positions of the ions.

18. Properties of Ionic Compounds • Another physical property of ionic compounds is their tendency to dissolve in water and conduct electricity while in solution. • Any compound that conducts electricity when dissolved in water is an electrolyte. In order to conduct electricity, ions must be free to move because they must take on or give up electrons.

19. Energy and Ionic Bonds • Endothermic- energy is absorbed during a reaction. • Exothermic- energy is released during a reaction. All ionic reactions of cations and anions are exothermic. Resultant compounds are more stable configuration (i.e. lower energy level) so excess energy is released. The amount of energy released is equal to amount needed to break the resultant bond.

20. Energy and Ionic Bonds • The energy required to separate one mole of the ions of an ionic compound is called lattice energy, which is expressed as a negative quantity. i.e. The greater (more negative) the lattice energy is, the stronger the force of attraction between the ions. Lattice energy tends to be greater for more-highly-charged ions and for small ions than for ions of lower charge or large size.

21. Ionic Compound Nomenclature • No single particle of an ionic compound exists so they are represented by a formula that provides the simplest ratio of the ions in an ionic compound and is called a formula unit. The overall charge of any formula unit is zero.

22. The charges of monatomic ions, or ions containing only one atom, can often be determined by referring to the periodic table or table of common ions based on group number. Ionic Compound Nomenclature

23. Ionic Compound Nomenclature • The charge of a monatomic ion is equal to its oxidation number. • The oxidation number, or oxidation state, of an ion in an ionic compound is numerically equal to the number of electrons that were transferred to or from an atom of the element in forming the compound.

24. Ionic Compound Nomenclature • In the formula for an ionic compound, the symbol of the cation is written before that of the anion. • Subscripts, or small numbers written to the lower right of the chemical symbols, show the numbers of ions of each type present in a formula unit.

25. Ionic Compound Nomenclature • If the ions in the ionic compound have the same charge, the formula unit contains one of each ion. • Na+ and Cl- combine to form NaCl. • Mg2+ and S2- combine to form MgS. • If the charges are not equal, we must balance the positive and negative charges. • Ca2+ and Cl- combine to form CaCl2. • Na+ and O2- combine to form Na2O.

26. Ionic Compound Nomenclature • In naming ionic compounds, name the cation first, then the anion. • Monatomic cations use the element name. • Monatomic anions use the root of the element name plus the suffix -ide. • If an element can have more than one oxidation number, use a Roman numeral in parentheses after the element name, for example, iron(II) to indicate the Fe2+ ion.

27. Crossover Rule • You can quickly verify that the chemical formula is written correctly by crossing over the charge on each ion. • The charge on the aluminum ion becomes the subscript for the oxygen, and the charge on the oxide ion becomes the subscript for the aluminum ion.

28. Compounds with Polyatomic Ions • Some ions contain more than one element. • An ion that has two or more different elements is called a polyatomic ion. • Although the individual atoms have no charge, the group as a whole has an overall charge.

29. Nick the Camel

30. Polyatomic Ions • Ionic compounds may contain: • Positive metal ions bonded to negative polyatomic ions, such as in NaOH sodium hydroxide • Negative nonmetal ions bonded to positive polyatomic ions, such as in NH4I ammonium iodide • Positive polyatomic ions bonded to negative polyatomic ions, such as in NH4NO3 ammonium nitrate.

31. Nomenclature with polyatomics • Follow the same rules as binary ionic compounds; if the charges are equal, the formula has one of each ion. • Mg2+ and SO42- combine to form MgSO4 Magnesium sulfate • K+ and ClO3- combine to form KClO3 Potassium Chlorate • If the charges are not equal, the total charge must equal zero. If you have more than one polyatomic ion, it is placed in parentheses. • Al3+ and CO32- combine to form Al2(CO3)2. Aluminum carbonate

32. Polyatomic Ions • To name a compound containing a polyatomic ion, follow the same rules as used in naming binary compounds. However, do not change the ending of the negative polyatomic ion name. i.e. To form a neutral compound, one calcium ion (Ca2+) must combine with one carbonate ion (CO3 2–) to give calcium carbonate with the formula CaCO3.

33. Ions of Transition Elements • Transition elements form positive ions just as other metals do, but most transition elements can form more than one type of positive ion and have multiple oxidation states. • Zinc and silver are two exceptions each forms one type of ion. Zinc ion is Zn2+ and the silver ion is Ag+. • A Roman numeral is used to indicate the oxidation number of a transition element ion.

35. Naming Ionic Compounds • Certain polyatomic ions, called oxyanions, contain oxygen and another element. • If two different oxyanions can be formed by an element, the suffix -ate is used for the oxyanion containing more oxygen atoms, and the suffix -ite for the oxyanion containing fewer oxygens.

36. Oxyanions • In the case of the oxyanions of the halogens, the following special rules are used. • four oxygens, per+root+ate (ex: perchlorate, ClO4–) • three oxygens, root + -ate (ex: chlorate, ClO3-) • two oxygens, root + -ite (ex: chlorite, ClO2-) • one oxygen, hypo+root+ite (ex: hypochlorite, ClO–)

37. Naming Ionic Compounds 1. NaBrO3 (sodium bromate) (magnesium nitrate) 2. Mg(NO3)2 (ammonium perchlorate) 3. NH4ClO4 4. Al(ClO)3 (aluminum hypochlorite)

38. Metallic Bonds and Properties of Metals • The bonding in metals is explained by the electron sea model, which proposes that the atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic cations. • These delocalized electrons are not held by any specific atom and can move easily throughout the solid. • A metallic bond is the attraction between these electrons and a metallic cation.

39. Properties of Metals • Metals generally have extremely high boiling points because it is difficult to pull metal atoms completely away from the group of cations and attracting electrons. • Metals are also malleable (able to be hammered into sheets) and ductile (able to be drawn into wire) because of the mobility of the particles. • The delocalized electrons make metals good conductors of electricity.