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Chapter 8: Ionic Compounds I. Forming Chemical Bonds

A.: force that holds two atoms together 1. Ionic bonds form through the attraction of a positive ion for a negative ion 2. Covalent bonds form through the attraction of a positive nucleus for negative electrons of another atom 3. Bonding behavior is determined by valence electrons i. Valence electrons are generally in s & p orbitals B. Ion Formation 1. Positive ions a. Formed when an atom loses electrons to reach a noble gas configuration b. Cation: positively charged ion c. Example: sodium ● 1s22s22p63s1 ● Loses 3s1 electron to get to noble gas configuration that it’s closest to (neon) ● Sodium ion: 1s22s22p6 ● Notation: Na+ d. Example 2: aluminum ● 1s22s22p63s23p1 ● Loses 3s2 and 3p1 electrons to get to noble gas configuration that it’s closest to (neon) ● Aluminum ion: 1s22s22p6 ● Notation: Al3+ 2. Negative ions a. Formed when an atom gains electrons to reach a noble gas configuration b. Anion: negatively charged ion c. Example: ● 1s22s22p63s23p5 ● Gains an electron in the 3p orbital to get to noble gas configuration that it’s closest to (argon) ● Ion: 1s22s22p63s23p6 ● Notation: Cl- d. Example 2: phosphorus ● 1s22s22p63s23p3 ● Gains 3 electrons in the 3p orbital to get to noble gas configuration that it’s closest to (argon) ● Ion: 1s22s22p63s23p6 ● Notation: P3- 3. Ions of Transition Metals a. Usually have an outer level of s2, then fill d orbital of second highest energy level b. Most transition metals lose the 2 electrons in their outer s orbital, forming ions with a 2+ charge c. Sometimes, the d orbital gets involved, forming ions of 3+ or greater d. All commonly formed ions of each element are listed to the upper right of the element symbol on your periodic table  Homework: p. 214 #1-5 II.Formation & Nature of Ionic Bonds A.Ionic bond: electrostatic force that holds oppositely charged particles (ions) together 1. Oxides: ionic compounds composed of a metal ion & oxygen 2. Salts: ionic compounds not containing oxygen 3. Binary compounds: contain only 2 elements 4. Example: sodium & chlorine ● Sodium (Na+) gives up its valence electron to chlorine (Cl-) ● Positive sodium is attracted to the negative chlorine ● Forms (NaCl) 5. Example 2: calcium & fluorine ● Calcium (Ca2+) gives up 2 valence electrons ● Fluorine (F-) only needs 1 electron, but the number of electrons lost must equal the number of electrons gained ● So… calcium gives one electron each to 2 fluorine atoms ● Forms calcium fluoride (CaF2) B. Properties of Ionic Compounds 1. Depend upon the bonds between the atoms involved 2. General characteristics: a. Form a crystal lattice: three-dimensional arrangement of particles; positive ions surrounded by negative ions & negative ions surrounded by positive ions i. Stronger attraction, harder & stronger crystal b. High melting & boiling points due to strong attraction between particles i. Stronger attraction means more energy is required to break the particles apart -- therefore, more heat required c. Soluble: dissolve in water i. When dissolved, these compounds become , meaning they conduct electricity C. Energy of Ionic Bonds 1. Chemical reactions always release or absorb energy 2. The formation of ionic compounds always releases energy, making it exothermic 3. Lattice energy: energy required to separate one mole of ions of an a. More negative lattice energy = stronger attraction b. Smaller ions have a more negative lattice energy than larger ions c. Bonds between ions with larger numerical charges have a more negative lattice energy Homework: ➔p. 217 #7-11 ➔p. 220 #12-17 III.Names & Formulas of Ionic Compounds A.Formulas for Ionic Compounds 1. Formula unit: simplest ratio of ions represented in an ionic compound a. Examples: NaCl, KBr, MgCl2 2. Ion: charged atom or a group of atoms with a charge a. Monatomic ion: one-atom ion with a definite charge i. Examples: Na+, Mg2+, O2-, F- ii. Charge is equal to an atom’s oxidation number: the number of electrons transferred from an element to form an ion (ex: Na+ has an oxidation number of +1) b. Polyatomic ion: group of atoms bonded together which have a net charge + - 2- i. Examples: NH4 , CN , SO4 B. Writing Formulas of Ionic Compounds 1. Rule: the sum of the charges in a stable compound must be zero a. Examples: i. Na+ + Cl- → NaCl + - 3 ii. Ag + NO3 → AgNO iii.Al3+ + Cl- → __(3) + __(-1) = 0 __(3) + 3(-1) = 0 AlCl3 (aluminum chloride) 2. How to write a formula: a. Step 1: find the chemical symbols or polyatomic ions & the charges they have based on the compound’s name (textbook, PT, common ions list) b. Step 2: write the charges in an algebraic formula, setting the sum of the charges equal to zero c. Step 3: determine how many of each ion are necessary to make the total charge equal to zero 3. Examples!! a. magnesium oxide d. ammonium sulfate 2+ 2- + 2- i. Mg O i. NH4 SO4 ii. __(2) + __(-2) = 0 ii. __(1) + __(-2) = 0 iii.MgO iii.2(1) + __(-2) = 0

b. calcium chloride iv.(NH4)2SO4 i. Ca2+ Cl- e. aluminum sulfate ii. __(2) + __(-1) = 0 i. Al3+ 2- iii.__(2) + 2(-1) = 0 SO4 iv.CaCl2 ii. __(3) + __(-2) = 0 c. aluminum sulfide iii.2(3) + 3(-2) = 0 3+ 2- i. Al S iv.Al2(SO4)3 ii. __(3) + __(-2) = 0 iii.2(3) + 3(-2) = 0

iv.Al2 S3 Homework: ➔p. 224 #19-23 (write out all 3 steps) ➔p. 225 #24-28 (write out all 3 steps) C. Naming Ions & Ionic Compounds 1. Naming Ions a. If two polyatomic ions are identical except for the number of oxygen atoms, the ion with more oxygens takes the suffix “-ate” and the ion with less oxygen takes the suffix “-ite” - - i. Examples: NO3 : nitrate, NO2 : nitrite 2- 3- SO4 : sulfate, SO4 : sulfite b. If four polyatomic ions are identical except for the number of oxygen atoms, use the prefix “per-” for the ion with the most oxygens, suffix “- ate” for the second most, suffix “-ite” for the third, and prefix “hypo-” with suffix “-ite” for the least - i. Examples: ClO4 : perchlorate - ClO3 : chlorate - ClO2 : chlorite ClO-: hypochlorite 2. Naming Binary (Ionic) Compounds a. Written from left to right on the periodic table b. Positive ion first -- use the full element name c. Negative ion last -- if it’s just one element (example: bromine), use the root of the element name and suffix “- ide” (bromide) d. If there is a polyatomic ion involved, just use the name of the polyatomic ion - - do not change the suffix to “-ide” + - i. Example: Na + NO3 → sodium + nitrate → sodium nitrate e. Transition metals & some metals on the right side of the table have more than one possible charge; to distinguish between these, use Roman numerals after the name of the element to indicate the charge i. Example: FeO → since there are no subscripts, they must have the same numerical charge. Oxygen is always 2-, so iron in this case has a charge of 2+. Name would be iron (II) oxide ii. Example: Fe2O3 → the subscript on one element is the same as the charge of the other. Oxygen’s charge is 2-, and iron in this case is 3+. Name would be iron (III) oxide

Homework: ➔copy figure 8-8 (p. 227) ➔p. 226 #29-33 ➔p. 227 #34-39 IV.Metallic Bonds & Properties of Metals A.Metals are not ionic compounds, but they do have some similar properties 1. Similarities are due to the fact that bonds are based on attractions between opposite charges B.Metallic Bonds 1. Form lattices in solid state, similar to crystal lattices 2. Metals’ outer electrons are loosely held & overlap when bonded a. Electron sea model: all of the valence electrons of the metals contained in a bond group together to form a “sea”; positive nuclei are “floating” in this b. These valence electrons are free to move from atom to atom -- called delocalized electrons c. Contributes to strong conductivity of metals C. Metallic Properties 1. Properties of metals depend upon # of valence electrons a. Group 1: one valence electron, soft b. Group 2: two valence electrons, harder 2. More valence electrons = greater hardness & strength a. Strongest bonds are in structural metals like nickel, chromium, & iron -- tend to share more electrons b. Transition metals usually form hardest & strongest materials 3. Properties of metallic bonds include high ductility, conductivity, malleability, durability D. Metal Alloys 1. : different metals combined to form one even stronger material 2. Two types of alloys: a. Interstitial i. Small holes in a metallic crystal are filled with smaller atoms ii. Example: carbon steel (98.5-99.95% Fe, 0.05-1.5% C); carbon fills the holes between atoms b. Substitutional i. Atoms of original metallic solid are replaced by other metal atoms of similar size ii. Example: brass (67-90% Cu, 10-33% Zn) Homework: ➔copy table 8-8 (p. 231) ➔p. 231 #40-44 ➔ionic compounds worksheet (activity Monday, quiz Tuesday) Chapter 8 Review: ➔p. 235: Vocabulary (all - 15 words) ➔p. 236: Concept Mapping 46 (DRAW IT) ➔p. 236: Mastering Concepts 47-59 ➔p .236: Mastering Problems 60-84 ➔p. 237: Mixed Review 85-96 ➔HONORS ONLY p. 238: Thinking Critically #97-104 ➔p. 239: Standardized Test Practice 1-10