Introduction to Bonding

1. Introduction to Bonding Ionic, Metallic & Covalent

2. Valence electrons • are the electrons in the highest occupied energy level (outer-most energy level) of an element’s atoms. • Crack out the Periodic Table

3. Electron Dot Diagrams • are diagrams that show valence electrons as dots • referred to as “Lewis Structures” named after the scientist that first used them • Show only Valence Electrons

4. Lewis Dot Examples

5. Octet Rule • Gilbert Lewis (1916) explained why atoms form certain kinds of ions and molecules. • In forming compounds, atoms tend to achieve the electron configuration of a noble gas(i.e. 8 valence electrons) • Atoms of the metallic elements tend to lose their valence electrons, leaving a complete octet in the next lowest energy level. • Atoms of some nonmetallic elements tend to gain electrons or to share electrons with another nonmetallic element to achieve a complete octet.

6. Formation of Ions • When an atom gains or loses electrons it becomes either positively or negatively charged • Cations are positively charged ions • Anions are negatively charged ions

7. Cations • Become POSTIVELY charged because they LOSE an electron • Metals have fewer than 4 valence electrons so they lose them to become cations. • 1+ means one electron was lost • 2+ means 2 electrons were lost • 3+ means 3 were lost

8. Transition Metal Cations • Fortransition metals, the charges of cations may vary. • An atom of Iron, for example, may lose two or three electrons. • the Fe2+ ion is called Iron (II) ion or Ferrous ion • the Fe3+ ion is called Iron (III) ion or Ferricion • This Roman numeral system is called the stock system.

9. Anions • Become NEGATIVELY charged because they GAIN electrons • Nonmetals have MORE than 4 valence electrons so they gain them to become anions. • 1- means one electron was gained • 2- means 2 electrons were gained • 3- means 3 were gained

10. Metalloids • Have 3, 4, or 5 valence electrons • Sometimes they gain electrons to become anions • Sometimes they lose to become cations • CARBON usually does not Ionize by itself, can form polyatomic ions. • Reference Table E, p. 2

11. Ionic compounds • When a positively charged cation is attracted to a negatively charged anion they form a relationship because opposites attract.

12. Example, table salt or sodium chloride • Na has one valence electron, which it loses easily (think the exploding water) • Chlorine wants one because it have seven already (deadly gas, react with the water in your lungs, don’t breathe it).

13. Na Cl ONE VALENCE ELECTRON SEVEN VALENCE ELECTRONS

14. Na Cl ZERO VALENCE ELECTRONS EIGHT VALENCE ELECTRONS

15. Na Cl

16. Salt (chemistry definition) • an ionic compound made between a metal and a nonmetal. • Although they are composed of ions, ionic compounds are electrically neutral. • The total positive charge of the cations equals the total negative charge of the anions.

17. Properties of Ionic Compounds • very stable structure. • high melting points. • (NaClhas a melting point of about 800°C.) • Are crystalline solids at room temperature. • (Crystal - a solid in which the atoms, ions, or molecules are arranged in repeating, three-dimensional patterns called a crystal lattice) • can conduct an electric current when dissolved in water.

18. Properties of Ionic Compounds • Water + Ionic Compounds • Ionic compounds also conduct electricity if they are dissolved in water. • When dissolved, the ions are free to move about in the aqueous (water) solution. • Solvation - a process that occurs when an ionic solute dissolves; in solution, solvent molecules surround the positive and negative ions.

19. Metallic Bonds • Metalsare best thought of as positive nuclei floating in a “sea of electrons”. • Thevalence electrons can drift freely from one part of the metal to another. • Metallic bonds consist of the attraction of the free-floating valence electrons for the positively charged nuclei

20. Metallic Properties • Metals are good conductors of electrical current because electrons can flow freely in them. • Metals are ductile—that is, they can be drawn into wires. • Metals are also malleable, which means that they can be hammered or forced into shapes.

21. Alloys • Alloysare mixtures composed of two or more elements, at least one of which is a metal. • Most of the metals you encounter are alloys. • Brass is an alloy of copper and zinc. • Sterling silver (92.5% silver and 7.5% copper) is harder and more durable than pure silver but still soft enough to be made into jewelry and tableware. • Bronze is an alloy of seven parts of copper to one part of tin. Bronze is harder than copper and more easily cast. • Non-iron alloys, such as bronze, copper-nickel, and aluminum alloys, are commonly used to make coins.

22. Amalgam • Liquid mercury has the ability to dissolve other metals to produce alloys and these alloys are called amalgams. • Amalgam examples are gold/mercury, silver/mercury and copper/mercury all of which are used in dentistry.

23. Covalent BondsAtoms held together by sharing a pair of electronsdiatomic molecule - a molecule consisting of two atoms e.g. • O2 • H2 • N2 • F2 • Br2 • I2 • Cl2

24. Covalent Continued • Water (H2O) is a molecule containing three atoms with two covalent bonds. • Two hydrogen atoms share electrons with one oxygen atom. • The hydrogen and oxygen atoms attain OCTETS by sharing electrons. • In the electron dot structures, the oxygen atom in water has two unshared pairs of valence electrons.