Chapter 20: CHEMICAL BONDING

1. Chapter 20: CHEMICAL BONDING

2. Classifications of Matter (p454) Matter Yes Yes Can it be separated by physical means? No No Mixture Pure substance Is the composition uniform? Can it be decomposed by ordinary chemical means? Yes No Yes No Homogeneous mixture Heterogeneous mixture Compound Element

3. Copper sulfate, CuSO4 is composed of • 1 copper (Cu) atom• 1 sulfur (S) atom, and• 4 oxygen (O) atoms. Chemical compounds (Page 602) • The atoms in compounds are chemically bonded and may only be broken down through chemical reactions. • A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms. Chemical bonds are electrical bonds

4. Elements and Compounds • Compounds formed when elements combine often have properties that are different from the properties of the elements from which the compound is formed. • Copper Sulfate – greenish coating on the Statue of Liberty • Copper- shiny colored metal • Sulfur- yellow powder • Oxygen- clear, odorless gas

5. Sodium and Chlorine • Na Sodium- a shiny soft, silvery metal that reacts violently with water. • Cl Chlorine – a poisonous greenish-yellow gas • NaCl Sodium Chloride- white crystals of ordinary table salt.

6. Chemical Formula (p603)

7. Why do elements combine to form compounds? • Elements combine in ways that produce the most stable electron arrangement. • The most stable electron arrangements have the outer energy shells filled. • For elements with 1 or 2 electrons the filled outer shell has 2 electrons. Hydrogen and Helium (Helium is stable). • For elements with > 2 electrons the filled outer shell has 8 electrons. Noble gases are all filled.

8. Element Structure • This neutral lithium atom has three positively charged protons, three negatively charged electrons, and four neutral neutrons.

9. How do elements combine to form compounds? • Atoms of elements gain, lose or share electrons in order to form a stable electron arrangement of 8 (or 2) electrons in the outer energy level. • Whether an atom gains, loses or shares electrons depends on how strongly the atom attracts electrons.

10. How strongly does an atom attract electrons? • The closer a negatively charged electron is to the positively charged nucleus, the more strongly it is attracted to the nucleus. • Moving down a group the outer level is farther from the nucleus • The more protons in the nucleus of the atom, the stronger the attraction for the electron. • Moving across a period the number of protons increases but distance stays the same. • Close is more important than number of Protons.

11. Periodic Table and Energy Levels • Look at the horizontal rows, or periods, in the portion of the table shown. • You can determine the number of electrons in an atom by looking at the atomic number written above each element symbol. Group number allows us to determine number of outer electrons. Increasing attraction for electrons  Decreasing attraction for electrons --

12. Noble Gases • Neon and the elements below it in Group 18 have eight electrons in their outer energy levels. • Their energy levels are stable, so they do not combine easily with other elements.

13. Halogens • The elements in Group 17 are called the halogens. • Fluorine is the most reactive of the halogens because its outer energy level is closest to the nucleus. • Fluorine attracts electrons more than any other element.

14. Halogens • Note, that if Fluorine gained an electron it would have the same electron arrangement as Neon.

15. Alkali Metals • Alkali metals each have one outer energy level electron. • It is this electron that is removed when alkali metals react. • The easier it is to remove an electron, the more reactive the atom is. • Unlike halogens, the reactivities of alkali metals increase down the group because the electrons are farther from the nucleus..

16. The Octet Rule • The octet rule states that chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons has an octet, or 8 electrons, in its outer shell.

17. Bonding Electrons are divided between inner shell and outershell electrons. Bonds form by the interaction of outer shell or VALENCE ELECTRONS.

18. Valence Electrons 18 1 IA VIIIA 2 13 14 15 16 17 IIA IIIA IVA VA VIA VII7A 9Ne For group 1 and 2 the number of valence electrons is equal to the group number. For group 13-18, the number of valence electrons is equal to the group number minus 10.

19. Using Dot Diagrams • Now that you know how to write electron dot diagrams for elements, you can use them to show how atoms bond with each other. • A chemical bond is the force that holds two atoms together. • Atoms bond with other atoms in such a way that each atom becomes more stable. That is, their outer energy levels will resemble those of the noble gases. 8 electrons

20. Types of Chemical Bonds (P608) • There are two extreme forms of connecting or bonding atoms: • Ionic—complete transfer of electrons from one atom to another. One atom gains and the other loses. • An atom that is no longer neutral because it has lost or gained an electron is called an ion (I ahn). • Covalent—sharing of electrons between atoms • Most bonds are somewhere in between.

21. Ionic Bonds—Loss and Gain (P609) • Sodium has only one electron in its outer level. • Removing this electron empties this level and leaves the completed level below with 8 electrons which is a stable configuration. • By removing one electron, sodium’s electron configuration becomes the same as that of the stable noble gas neon. • Removing this electron also changes the sodium atom into a positively charged ion. Na  Na+

22. Ionic Bonds—Loss and Gain • Chlorine forms bonds in a way that is the opposite of sodium—it gains an electron. • When chlorine accepts an electron, its electron configuration becomes the same as that of the noble gas argon and the atom is changed into a negatively charged ion. - Cl Cl 

23. Bond Formation • The positive sodium ion and the negative chloride ion are strongly attracted to each other. • This attraction, which holds the ions close together, is a type of chemical bond called an ionic bond. • A compound is a pure substance containing two or more elements that are chemically bonded.

24. Ionic Bonds Essentially complete electron transfer from an element of low attraction (metal) to an element of high electron attraction (nonmetal) 2 Na(s)+ Cl2(g) 2 Na+ + 2 Cl- Therefore, ionic compounds exist primarily between metals at the left of the periodic table (Groups 1, 2, and transition metals) that form + ions and nonmetals at the right (Groups 16-17) that form - ions. Cl- Na+

25. Lose 2 Gain 2 • Some elements like magnesium, Mg, in Group 2 has two electrons in its outer energy level. • Magnesium can lose two electrons and achieve a completed energy level. • Some atoms, such as oxygen (group 16), need to gain two electrons to achieve stability. • The two electrons released by one magnesium atom could be gained by a single atom of oxygen. Magnesium Oxide is formed.

26. Lose 2 Gain 2 (group 2 and 16) Magnesium and Oxygen • When this happens, magnesium oxide (MgO) is formed. • Mg Mg+2 +2e- and O+ 2e-  O-2 • Mg+2 + O-2 ionic bond • 2Mg +O2 2MgO

27. Magnesium and Chlorine Lose 2 Gain 1 (group 2 and 17) • Mg+2 + 2e- + 2 Cl Mg+2 +2Cl-1 MgCl2 Magnesium Chloride

28. Lose 1 Gain 2 (which groups) Na+ O-2 Na2O sodium oxide

29. Metallic Bonding—Pooling • In a metal, the electrons in the outer energy levels of the atoms are not held tightly to individual atoms. • Instead, they move freely among all the ions in the metal, forming a shared pool of electrons. • The outer electrons in metal atoms readily move from one atom to the next to transmit current.

30. Metallic Bonding—Pooling • Metallic bonds form when metal atoms share their pooled electrons. This bonding affects the properties of metals.

31. Covalent Bonds—Sharing • Some atoms are unlikely to lose or gain electrons because the number of electrons in their outer levels makes this difficult. • The alternative is sharing electrons. • Many nonmetal elements form covalent bonds with other non metals. • These bonds follow the octet rule. • The chemical bond that forms between nonmetal atoms when they share electrons is called a covalent(koh VAY luhnt) bond.

32. The Covalent Bond • You can see how molecules form by sharing electrons equally in this figure.

33. Electron Dot Diagrams of Electron Sharing Cl + Cl Cl Cl2  Cl

34. Forms of Chemical Bonds • There are two extreme forms of bonding atoms. Most bonds are somewhere in between. • ionic • covalent • Polar Covalent Bond — the electrons are not shared equally between atoms • Read(P612-614) • Nonpolar Covalent Bond- • electrons are shared equally between atoms d- d+ d+

35. Diatomic GasesNon Polar Bonds (equal sharing) • Br2- bromine • I2 – iodine • N2 – nitrogen • Cl2 - chlorine • H2 – hydrogen • O2 – oxygen • F2 – fluorine • Mr. BrINClHOF

36. Polar Bonds and Polar Molecules • Some atoms of some elements have a greater attraction for electrons than others do. • A polar bond is a bond in which electrons are shared unequally.

37. Chapter 20 Section 3Writing Formulas and Naming Compounds Page 615

38. Chemical Formulas • A chemical formula is a combination of chemical symbols and numbers (subscripts) that show which elements are present in a compound and how many atoms of each element are present. • When no subscript is shown, the number of atoms is understood to be one. • A water molecule contains one oxygen atom and two hydrogen atoms, so its formula is H2O. • Binary Compounds- a compound composed of two elements.

39. Oxidation Number (P616) • Oxidation number is the number of electrons lost, gained or shared by an element in a compound. • The oxidation number can be determined from the periodic table. • For Ionic compounds (ionic bonds) the oxidation number is the charge on the ion. • Sometimes we are given the oxidation number (Table 2) for transition metals.

40. Oxidation Number (p616)Oxidation number is the number of electrons lost, gained or shared by an element in a compound. 1+ 2+ 3+ 4+ 3- 2- 1- 0

41. Special Ions (Table 2 P 616)

42. Subscript/Superscript • Subscript written below the symbol shows how many atoms or ions are present in the compound. • Superscript written above the symbol shows the charge or oxidation number of the atom or ion and is used in knowing how to write the formula • but is not actually written in the final formula.

43. Writing Chemical Formula (P617) • Write the symbol of the element or polyatomic ion that has a positive oxidation number or charge. Hydrogen, the ammonium ion (NH41+), and all metals have positive oxidation numbers. • Write the symbol of the element or polyatomic ion with the negative oxidation number. Nonmetals other than hydrogen and most polyatomic ions have negative oxidation numbers • The charge without the sign of one ion becomes the subscript of the other ion. Reduce the subscripts to the smallest common multiple. • Mg2+ Cl1-  MgCl2

44. Example Writing Formula • Aluminum Oxide Al O Al O • Magnesium Sulfide Mg S Mg S • Lithium Oxide Li O Li O • Potassium Sulfide K S K S • Calcium Nitride Ca N Ca N Applying Math Page 617: 1,2

46. Write the formula for the following ionic compounds: • sodium iodide • copper(II)sulfate • calcium chloride • sodium carbonate • potassium sulfide • calcium nitrate • lithium nitrate

47. Writing Names –Ionic Bonds(binary compounds) P618 • Write the name of the positive ion. • Some positive ions may have more than one oxidation number. (see table 2) Use the Roman numeral to identify the ion charge. Negative ions have only one oxidation number. • Write the root name of the negative ion. Drop the last syllable • Add the ending –ide.

48. Determining the type of Bond • Metallic Bonds- between transition metals and other transition metals • Ionic Bonds- between elements on one side of the periodic table with an element on the other side. Non-metals bonding with metals. • Covalent Bonds- between non-metals, elements close together on the right side of the PT.

49. Examples Writing Names • BaF2 • NaCl • K2S • MgCl2

50. Element Names in Binary Compounds