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IONIC BONDS

IONIC BONDS. Sec. 1: Forming Chemical Bonds. How do the thousands of compounds form from the relatively few elements known to exist? The answer is found in the electron structure of the atoms of the elements involved and the nature of the forces between these atoms.

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IONIC BONDS

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  1. IONIC BONDS

  2. Sec. 1: Forming Chemical Bonds • How do the thousands of compounds form from the relatively few elements known to exist? • The answer is found in the electron structure of the atoms of the elements involved and the nature of the forces between these atoms. • CHEMICAL BOND: the force that holds two atoms together

  3. Do you remember……… • Elements within a group have similar properties? • Each of the groups of representative elements have the same number of valence electrons? • All elements (except for H, He, Li, Be) want to have 8 electrons in their outer shell? This is called the what? • OCTET RULE

  4. Valence Electrons

  5. And do you remember this? • Electron-dot structures illustrate the valence electrons? • Ionization energy refers to how easily an atom loses an electron? • Electronegativity is the tendency to want to gain electrons? • The difference in reactivity is directly related to the valence electrons? • Elements tend to react to acquire the stable electron structure of a noble gas

  6. Electron Configuration • Let’s look at the electron configurations of POTASSIUM and ARGON. (I will put them on the chalkboard.) • Why is argon stable? • How can potassium achieve a noble gas configuration? • What is the electron configuration for a potassium ion? • What is the symbol notation for the potassium ion?

  7. Formation of Positive Ions • When metals lose electrons, a POSITIVE ION is formed. • A positively charged ion is called a CATION. So, metals are called cations. • Reactivity of metals is based on the ease with which they lose valence electrons to achieve a stable octet, or noble gas configuration. • When elements in groups 1B – 4A in periods 4 through 6 lose electrons so the outer energy level contains full s, p, and d sublevels, they are relatively stable. They do not have a noble gas configuration, but they are pretty stable. We call this pseudo-noble gas configurations.

  8. Formation of Negative Ions • When nonmetals gain electrons, a NEGATIVE ION is formed. • We call these negatively charged ions ANIONS. Nonmetals are called anions. • To designate an anion when naming compounds, we add the ending –ide to the root name of the element. • EX: when chlorine gains an electron and becomes an anion, we call it chloride. • What would the rest of the nonmetals be named? (Let’s look at the periodic table.)

  9. Forming Negative Ions, continued: • Nonmetals gain the number of electrons that, when added to their valence electrons, equals eight.

  10. Section 1 Assessment 1. What is a chemical bond? 2. Why do ions form? 3. What family of elements is relatively unreactive and why? • (a.) Describe the formation of a positive ion. (b.) Describe the formation of a negative ion.

  11. 5. Predict the change that must occur in the electron configuration if each of the following atoms is to achieve a noble gas configuration: (tell me if they gain or lose electrons AND how many) • Nitrogen • Sulfur • Barium • Lithium

  12. Sec. 2: The Formation and Nature of Ionic Bonds

  13. *Sodium, a metal, LOSES an electron and becomes POSITIVELY charged. *Chlorine, a nonmetal, GAINS an electron and becomes NEGATIVELY charged. The compound, sodium chloride, forms because of the attraction between oppositely charged sodium and chloride ions. IONIC BOND: the electrostatic force that holds oppositely charged particles together in an ionic compound

  14. Compounds that contain ionic bonds are called ionic compounds. • Metals + oxygen = OXIDES • Metals + most other nonmetals = SALTS • BINARY IONIC COMPOUNDS: ionic compounds that contain ONLY two different elements: one metal (cation) and one nonmetal (anion)

  15. REMEMBER!!!!!!!!!!!!!!!!! • In a compound, the number of electrons lost MUST EQUAL the number of electrons gained! • The charge of the ionic compound is ZERO!!!!!!!!!!!!!!

  16. Let’s look at page 216. • There are 4 ways we can show the formation of an ionic compound: • Electron configuration • Orbital notation • Electron-dot structures • Atomic models We will look at these more tomorrow.

  17. Properties of Ionic Compounds • 1. Ionic compounds form CRYSTALS and a CRYSTAL LATTICE. Large numbers of (+) ions and (-) ions exist together in a packed, regular, repeating pattern that balances the forces of attraction and repulsion between the ions.Each positive ion is surrounded by negative ions and each negative ion is surrounded by positive ions.

  18. 2. High melting point • 3. High boiling point • 4. hard, rigid, and brittle • 5. formation of ionic compounds is ALWAYS exothermic. • 6. Do not conduct electricity in the SOLID state. But, they DO conduct electricity in the liquid state or when dissolved in water. • Why do you think this is so????? • An ionic compound whose aqueous solution conducts an electric current is called an electrolyte.

  19. Energy and the Ionic Bond • The energy required to separate the ions in an ionic compound is called the lattice energy. • It is also the energy given off when positive and negative ions attract. • Smaller ions usually have more lattice energy because the nucleus is closer to and thus has more attraction for the valence electrons.

  20. Ionic Radii • Back in ch. 6, we mentioned IONIC RADUIS. Two things we need to know: • The ionic radius becomes smaller for cations (lose electrons) as you go from left to right on the periodic table. • The ionic radius becomes larger for anions (gain electrons) as you go from left to right on the periodic table. • Look on page 166.

  21. Section 2 Assessment • 1. What is an ionic bond? • 2. How does an ionic bond form? • 3. List four physical properties associated with an ionic bond.

  22. 4. Describe the arrangement of ions in a crystal lattice. each ion is surrounded by oppositely charged ions

  23. Homework: Copy and fill in

  24. Sec. 3: Names and Formulas for Ionic Compounds • Chemical formulas and the names for compounds must be understood universally. Therefore, a set of rules is used in the naming of compounds.

  25. Formulas for Ionic Compounds • The simplest ratio of the ions represented in an ionic compound is called a formula unit. EXAMPLES OF FORMULA UNITS: • KBr = one K ion, and one Br ion • Na3P = three Na ions, and one P ion • Be2+ = beryllium ion (lost 2 electrons) • N3- = nitrogen ion (gained 3 electrons)

  26. Monatomic ion is a one-atom ion, such as Be2+ and N3- . It is made up of ONLY one element. • The charge of a monatomic ion is its OXIDATION NUMBER. (It represents the number of electrons gained or lost.) • Groups 1A, 2A, 5A, 6A, and 7A each have only one oxidation number. • Groups 3A and 4A metals and most transition metals have more than one oxidation number. We use Roman numbers—we will get back to this. • Look at the table of common ions on page 222. (This is a TOOL, like the periodic table. You will not have to memorize it, but you might want to make notes on your periodic table. Just keep up with it!)

  27. Oxidation Numbers….. • can also be called the oxidation state. • equals the number of electrons gained or lost by an element to form an ion. • used to determine the formulas for the ionic compounds they form.

  28. Naming Binary Ionic Compounds • Use the full name of the cation (metal). • Use the name of the anion (nonmetal) but change the ending to “-ide”. H = hydride F = fluoride N = nitride Cl = chloride P = phosphide Br = bromide O = oxide I = iodide S = sulfide

  29. Examples: What is the name of the following compounds? (a.) MgCl2 (b.) LiBr

  30. Quickly and quietly, you do these: (1.) K2O (2.) Mg3N2 (3.) Al2O3 (4.) Na3P

  31. What about compounds with transitional metals and Group 3A and 4A metals? • ROMAN NUMERALS are always used with transition metals when writing and naming formulas. Put them in parentheses. • The Roman numeral indicates the oxidation number. • When naming ionic compounds with transition metals, to know what the Roman numeral is, look at the anion (nonmetal) and assign it its ox. #. Then you can figure out the cation’s ox. #. LOOK AT THE EXAMPLES I DO ON THE BOARD.

  32. EXAMPLE: what is the name of CuCl2 ? Look at the ox. # of the anion (Cl) = -1 Multiply the ox.# by the subscript = (-1)(2)= -2 The Cu must have a +2 to cancel out the -2. Use Roman numbers to show the ox. # of Cu. The answer is copper (II) chloride.

  33. More examples: (1.) CdBr2 (2.) CrCl3 (3.) CuCl2 (4.) PbS2

  34. Writing Binary Ionic Compounds • Write the symbols for each element. • Assign each element their oxidation #. • SWAP ox. #’s and write as the subscript, and DROPthe charges. (NEVER write “1” as a subscript and never write the charge on the final formula!) Let’s look on the overhead for some examples from page 224 (the top).

  35. What about compounds with transitional metals and Group 3A and 4A metals? • We will use the same rules as for the other metals, EXCEPT to know what the charge is for the metal, we simply look at the Roman numeral – it IS the charge for that metal. Easy! • EXAMPLE: What is the formula for copper (II) chloride? Cu2+ Cl1- (swap and drop) CuCl2

  36. Try these on your own: (1.) iron (III) oxide (2.) palladium (II) chloride (3.) lead (IV) oxide** (4.) cobalt (II) sulfide**

  37. Section 4: Metallic Bonds and Properties of Metals • Although metal atoms have at least one valence electron, they do not share these electrons with neighboring atoms nor do they lose electrons to form ions. • Instead, the outer energy levels of the metal atoms overlap. The electron sea model proposes that all metal atoms in a metalic solid contribute their valence electrons to form a “sea” of electrons. (See pg. 228 – 229.) • This explains the properties of metallic solids such as malleability, conduction, and ductility.

  38. The electrons in the outer energy levels of the bonding metallic atoms move easily from one atom to the next and are called delocalized electrons. • A metallic bond is the attraction of a metallic cation for delocalized electrons.

  39. Properties of Metals • 1. Melting points vary greatly, but have moderately high melting points and high boiling points. • 2. Metals are malleable and ductile. The particles in metallic bonds can be pushed or pulled past each other. (See fig. 8-10.)

  40. As a metal is struck by a hammer, the atoms slide through the electron sea to new positions while continuing to maintain their connections to each other.

  41. 3. Metals are generally durable. They are strongly attracted to the electrons surrounding them and aren’t easily removed from the metal. • 4. Metals are good conductors of heat and electricity because of the mobile electrons. (Let’s look at the next to the last paragraph on pg. 229.)

  42. 5. Metals have luster. The delocalized electrons interact with light, absorbing and releasing photons. • 6. The more delocalized electrons, the harder and stronger they are because the “d” orbital electrons are also mobile.

  43. Metal Alloys • An alloy is a mixture of elements that has metallic properties. • Look at the table of some alloys on page 231. Which alloy would be used to make trophies? • The properties of alloys differ somewhat from the properties of the elements they contain. For example, steel is stronger than iron, but does have some properties of iron.

  44. Two Basic Types of Alloys • 1. Substitutional alloy: atoms of the original metallic solid are replaced by other metal atoms of similar size EX: sterling silver: Cu replaces Ag Other examples are brass, pewter, and 10-carat gold

  45. 2. Interstitial alloy: the small holes (interstices) in a metallic crystal are filled with smaller atoms. EX: carbon-steel: holes in the Fe crystal are filled with carbon atoms The presence of carbon makes the iron solid harder, stronger and less ductile than pure iron.

  46. ASSIGNMENT: pg. 236-237 #46, 47, 48, 52, 57, 60, 62, 63, 72, 74, 75, 76, 77, 80, 82

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