1. 1 Ionic and Metallic Bonding Ch. 7

2. 2 Review What is a valence electron? –Electrons in the highest (outermost) occupied energy level Related to the group number on the periodic table Determines chemical properties of an element

3. 3 What are electron dot structures? Dots around the symbol for the element to represent the valence electrons –Examples: H (Valence # 1)Examples

4. 4 What is the octet rule? Atoms want to achieve 8 electrons in outer shell Metals lose their valence electrons –Leaving a complete octet in the next lowest energy level Nonmetals tend to gain or share electrons in order to achieve 8 electrons in their outer shell Which group of elements already has a filled outer shell? –Nobel gases so no need to react

5. 5 Ions Ion – charged atom Atoms are no longer neutral when they gain or lose electrons Two types of ions 1.Anion – An ion with a negative (-) charge gained electrons Nonmetals form anions Anions are always larger than their corresponding neutral atom Name of anions end in ide 2.Cation – An ion with a positive (+) charge lost electrons Metals form cations Cations are always smaller than their corresponding neutral atom

6. 6 Example Fluorine has a valence # of 7 –it will gain 1 electron to fill its shell and reach 8 –  fluorine becomes a negative ion F 1  –For some atoms their ions have different names, Fluoride ion comes from fluorine. Sodium, valence # 1 –it will lose its electrons and become Na 1+ If your gaining an electron then why are you forming a negatively charged ion (anion)?

7. 7 How many valence electrons does the neutral sodium atom have? Explain why the sodium ion has a charge of 1+. Is this ion a cation or anion? What is the group number for sodium?

8. 8 How many valence electrons does the neutral magnesium atom have? Explain why the magnesium ion has a charge of 2+. Is this ion a cation or anion? What is the group number for magnesium?

9. 9 How many valence electrons does the neutral aluminum atom have? Explain why the Aluminum ion has a charge of 3+. Is this ion a cation or anion? What is the group number for Aluminum?

10. 10 How many valence electrons does the neutral Chlorine atom have? Explain why the chlorine ion has a charge of 1-. Is this ion a cation or anion? How many electrons will chlorine gain? What is the group number for chlorine?

11. 11 How many valence electrons does the neutral Sulfur atom have? Explain why the Sulfur ion has a charge of 2-. Is this ion a cation or anion? How many electrons will Sulfur gain? What is the group number for Sulfur?

12. 12 How many valence electrons does the neutral Phosphorus atom have? Explain why the Phosphorus ion has a charge of 3-. Is this ion a cation or anion? How many electrons will Phosphorus gain? What is the group number for Phosphorus?

13. 13 Gain or lose? Why do certain groups gain electrons while other groups lose electrons to form a stable octet? –Depends on the atomic radius –Smaller atomic radius harder it is to remove electrons so it is easier for them to gain electrons General rules: 1.If your element has < 4 valence electrons it will lose those electrons 2.If your element has > 4 electrons it will gain electrons How many electrons will the element gain? –You want to fill the outer shell

14. 14 Energy for Ion Formation Electron can move to higher energy level when the atom absorbs energy –Cations form when electrons gain enough energy to escape from atoms. Ionization - process of removing electrons and forming ions. Ionization Energy – amount of energy used to remove an electron –Lower the ionization energy the easier it is to remove an electron –Higher the ionization energy the harder it is to remove an electron Electron affinity - tendency of an atom to attract electrons

15. 15 What is Chemical Bonding? Chemical bonding – force that holds atoms or ions together; force of attraction between the atoms in a substance Atoms of elements are combined to form new substances Electron arrangement of the outermost energy level of an atom determines if atoms will form bonds. –Atoms want to complete their outermost energy level.

16. 16 Ionic Bonding Ionic bonding includes the transferring of electrons it occurs because of the attraction of positive and negative ions for each other Ion - a charged atom One atom gains electrons and the other loses electrons. Substances that are ionically bonded form ionic solids Ionic compounds are electrically neutral

17. 17 Properties of Ionic Compounds Properties of an ionic compound can be explained by the strong attractions among ions within a crystal lattice. Properties 1.High melting point 2.Poor conductors of electric current when solid 3.Ionic compounds can conduct an electric current when melted or dissolved in water 4.Most are crystalline solids at room temperature

18. 18 Ionic Bonding of NaCl After Bonding After Na gives 1 electron to Cl, It now has a filled outer energy level, But 11+ and 10-, a charge of +1 After Cl receives 1 electron from Na, it Now has a filled outer energy level, but 17+ and 18-, a charge of -1

19. 19 What are Chemical Formulas? Chemical Formulas are combinations of chemical symbols –A group of symbols that shows the makeup of a compound –a shorthand way of representing chemical compounds: H 2 O and sometimes a molecule of an element: H 2 Why are there two hydrogen atoms for every oxygen atom?

20. 20 Oxidation Number # of valence electrons determines how atoms will bond. Oxidation number - # of electrons an atom gains, loses, or shares when it forms chemical bonds. (describes combining capacity) Example: –Na 1+ (ion formed when Na loses 1 electron) 1+ is the oxidation number! –Cl 1  (ion formed when gains 1 electron) 1- is oxidation number

21. 21 Using Oxidation Numbers Can use oxidation # of atoms to predict how atoms will combine and what the formula for the resulting compound will be. –Follow this rule: the sum of the oxidation numbers of the atoms in a compound must be zero –Na 1+ + Cl 1  =NaCl (1-1=0)

22. 22 Polyatomic Ions A group of atoms that act as a single atom and that have a charge are called Polyatomic Ions –Although bonds with ion are covalent, these ions usually form ionic bonds –Examples: NH 4 1+, OH 1-

23. 23 Cross-Over Method 1.Write the symbol for the cation (+) first, then write the symbol for the anion (-). 2.Place the oxidation number above the ion’s symbol as a superscript. –Example: A +x B -y Superscript – small number written above and to the right of the symbols 3.Drop the positive and negative sign –Example: A x B y 4.Cross the number to the lower position of the opposite ion making a subscript –Example: A y B x

24. 24 Cross over continued If the charges are the same number, DO NOT write the subscripts –Example: AB (if x=1 and y=1) Do NOT change the subscripts of polyatomic ions Combined a polyatomic ion with another ion with a different charge that is greater than 1, place parenthesis around the polyatomic ion before placing the new subscript to the right of the parenthesis –Example: Al +3 SO 4 -2 Al 2 (SO 4 ) 3

25. 25 Naming Ionic Compounds Rule 1: (metals and nonmetal) –Write the 1 st element from the periodic table then the second element changing the ending to (ide) Example: KCl Potassium Chloride Aluminum Hydroxide Rule 2: (metal and polyatomic ion) - Write the 1st element from the periodic table then write the name of the polyatomic ion Al(OH) 3

26. 26 Naming Ionic compounds with Transitional metals Transitional metals can change their valence number –So will have more than one oxidation number When naming need to include the oxidation number of the transitional metal as a Roman numeral. Example –Fe 2 O 3 is Iron (III) oxide –FeO is Iron (II) oxide

27. 27 Metallic Bonds Metals give up electrons easily. –What happens when trying to bond two metals? Metallic bonds - outer electrons of the atom form a common electron cloud –electrons sort of become the property of all the atoms - “sea of electrons” Positive nuclei of atoms of metals are surrounded by mobile electrons that are all attracted by the nuclei at the same time. The “sea of electrons” is why metals are ductile, malleable, and good conductors

28. 28 Metallic Bonds Continued Metallic Bond – attraction between a metal cation and the shared electrons that surround it Strength of the metallic bond depends on the number of valence electrons –More valence electrons – stronger bond –  alkali metals have relatively weak bonds, are soft, and have relatively low melting points

29. 29 Properties of Metals Structure within a metal affects the properties of the metals Examples: –Ability to conduct electricity Since metals have an electron cloud it is easy for electrons to flow –Malleability Lattice of the metal is flexible so if hit it with a hammer you don’t shatter the metal you just move the ions around