1 / 37

Chapter 22: Chemical Bonding

Chapter 22: Chemical Bonding. Section 1– Electrons and Chemical Bonding. Chemical bonding – joining of atoms to form new substances Chemical bond – interaction that holds 2 atoms together. Valence Electrons. e - in outermost shell Determines an atom ’ s chemical properties

kirsi
Download Presentation

Chapter 22: Chemical Bonding

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 22: Chemical Bonding

  2. Section 1– Electrons and Chemical Bonding • Chemical bonding – joining of atoms to form new substances • Chemical bond – interaction that holds 2 atoms together

  3. Valence Electrons • e- in outermost shell • Determines an atom’s chemical properties • Used to form bonds • Within a group, or family, atoms have the same # of valence e- • Atoms with fewer than 8 valence e- are more likely to form bonds than an atom with 8 e-

  4. Types of Bonds • Ionic, Covalent, Metallic • Atoms bond by sharing, gaining or losing e- to have a filled outermost energy level. • A full set = 2 e- for a few of the elements • A full set = 8 e- for most elements • Which ones? Why?

  5. Section 2: Ionic Bonds • Bonds form by gaining or losing e-, resulting in charged atoms called ions. • Oppositely charged ions are attracted to one another • Metal + Nonmetal • Positive and negative charges cancel each other out to form an overall neutral compound

  6. Metal Atoms • Have few valence electrons • Usually lose these valence e- and form positive ions (cations) • Some transition metal ions can have multiple charges. For example, iron can have a +2 or a +3 charge. • The charge is written as a superscript of the symbol: Ex. K+, Ca2+, Al3+

  7. Nonmetal atoms • Have almost full valence shells • Tend to gain e- from other atoms and form negative ions (anions) • The charge is written as a superscript of the formula: • Ex. P3-, S2-, Cl-

  8. Polyatomic Ions • Poly = “many” • Polyatomic = “many atoms” • A group of atoms that behave as a single ion with an overall positive or negative charge • Treat the polyatomic ion as a single unit with a single charge.

  9. Writing formulas for ionic compounds • The number of positive charges and negative charges must balance in an ionic compound • The formula represents this balance • Subscripts are used to indicate the ratio of elements in the compound (no “1”)

  10. Writing formulas for ionic compounds • Find oxidation number (charge) for both parts • For elements in groups 1 and 2, use group #. Boron family is 3+, Carbon family ± 4, Nitrogen family is -3, Oxygen family is -2 and halogens are -1. • For cations followed by a roman numeral, the roman numeral is the oxidation # • For polyatomic ions check the list. Do not change the subscripts within the polyatomic ion formula.

  11. 2. Write symbols. The positive ion first and negative ion second. 3. Put polyatomic ions in parentheses if more than one is needed. 4. Use subscripts to designate the number of each part for the total + and – charges to be = a. Find the least common multiple of both charges b. determine the factor needed to get that charge and use that as the subscript.

  12. Sodium sulfide • Potassium iodide • Lithium oxide • Barium fluoride • Iron(III) oxide • Copper(II) chloride g. Sodium acetate h. Zinc(II) carbonate i. Chromium(II) sulfate j. Cobalt(III) iodide

  13. Answers • Na2S • KI • Li2O • BaF2 • Fe2O3 f. CuCl2 g. NaC2H3O2 h. ZnCO3 i. CrSO4 j. CoI3

  14. Writing names for ionic compounds • Write the names of the + and - part of the formula • The + part is the name of the element or polyatomic ion • Check the list of elements to see if it needs a roman numeral. If so, use the negative part of the formula to figure out the positive charge on the metal. • To name the second/- part, if it is an element change the ending to “-ide” • If it is a polyatomic ion, keep the name as is

  15. Practice LiCl MgCl2 BeO CaCl2 HgS SnF2 g. (NH4)3PO4 h. ZnCO3 i. Sn(OH)2 j. Li2SO4 k. KC2H3O2

  16. Answers • Lithium chloride • Magnesium Chloride • Beryllium oxide • Calcium Chloride • Mercury(II) sulfide • Tin(II) fluoride g. Ammonium phosphate h. Zinc(II) carbonate i. Tin(II) hydroxide j. Lithium sulfate k. Potassium acetate

  17. Section 3: Covalent and Metallic Bonds • Covalent bonds are formed when atoms share one or more pairs of valence e-. • Forms between 2 nonmetals. • Covalent bonds result in the formation of molecules. • Define molecule

  18. Octet Rule • Atoms combine in such a way so as to fill the valence shell (usually that means 8 e- but could be just 2 e-)

  19. How many bonds? • The number of e-that an atom needs to fulfill it’s valence is equal to the number of covalent bonds it can form. • Ex. N can make 3 covalent bonds because it has 5 e- in its valence shell • Ex. H can make 1 covalent bond because it has 1 e- its valence shell.

  20. Diatomic Elements • Certain elements exist as pairs in nature because that is how they are most stable. • Di = 2 • Just remember Professor BrINClHOF (Bromine, Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen and Fluorine) • You need to memorize the 7 diatomic elements!!

  21. Practice! Draw a Bohr diagram and Lewis structure for the following: • Water (H2O) • Diatomic Fluorine (F2) • Silicon tetrafluoride (SiF4)

  22. Double Bonds • When atoms share 2 pairs of e-, it is a double bond

  23. Triple Bonds • When atoms share 3 pairs of e-, it is a triple bond

  24. Naming Covalent Compounds • Many compounds have common names such as "methane", "ammonia" and "water” • Simple covalent compounds can be named using prefixes to indicate how many atoms of each element are in the formula. • The ending of the last (most negative) element is changed to -ide.

  25. Naming Covalent Compounds Prefixes: • Mono = 1 • Di = 2 • Tri = 3 • Tetra = 4 • Penta = 5 • Hexa = 6 *If there is just 1 of the first element no prefix is used

  26. Practice • CO2 • SiF4 • CO • NBr3 • P2O5 • BCl3

  27. Answers • CO2 – carbon dioxide • SiF4 – silicon tetrafluoride • CO – carbon monoxide • NBr3 – nitrogen tribromide • P2O5 – diphosphorus pentoxide • BCl3 - boron trichloride

  28. Practice • Carbon tetrabromide • Phosphorus triiodide • Bromine • silicon monoxide • Silicon disulfide

  29. Answers • CBr4 • PI3 • Br2 (Bromine is a diatomic molecule!) • SiO • SiS2

  30. Metallic Bonds • A bond formed by the attraction between positively charged metal ions and surrounding e- • Think of a metal as being made up of positive ions with electrons “swimming” around keeping the ions together

  31. Properties of Metals • Because e- can move freely about, metals have particular propeties.

  32. Conduct electricity well

  33. Metals can be reshaped (ductile, malleable)

  34. Metals can bend without breaking

More Related