Chapter 15 Notes: Ionic Bonding and Ionic Compounds

Chapter 15 Notes: Ionic Bonding and Ionic Compounds Reading Assignment C15! 1. Read C15 pp. 412-435 and while reading, continue building your science vocabulary table that includes all terms in bold face type and all terms you are unfamiliar with or unsure of. Your vocabulary table needs to include three things: 1. The term 2. The definition 3. A picture of what the term means to you. 2. On page 431, organize the Concept Map into a form that correctly connects the major ideas of the chapter. 3. Read C15 pp 413-425. Answer section review questions 1-14 and write a list of 5 precious gems (not diamonds) and/or minerals including their chemical formulas. Example. Sapphires belong to the mineral class corundum. The formula is Be3Al2SiO6

I. Atoms and Ions a. Atoms are electrically neutral- equal number of electrons and protons. b. Ion- atoms which have either gained or lost electrons- contain a net charge due to imbalance between protons and electrons. i. Cation- electrically positive ion- due to a loss of electrons. ii. Anion- electrically negative ion- due to a gain of electrons. c. Valence shell - Highest energy level occupied by electrons. i. Valence electrons - electrons in the valence shell. 1. Elements in the same groups usually have the same number of valence electrons. 2. The group number for representative elements is the same as the number valence electrons. d. Electron dot structures or Lewis-Dot structures (G.N. Lewis) i. Lewis, G.N., Published The Atom and the Molecule in the Journal of the American Chemical Society, 1916. 1. Describes the valence electrons in an atom utilizing dots and symbols (x can be used to represent valence electrons, especially when representing those involved in chemical bonds) 2. Drawing Lewis Structures C16.1 p.441. e. Octet Rule- Atoms have a tendency to gain/lose electrons and achieve electron configurations of the noble gases. i. First proposed by Irving Langmuir in 1919. Published in The structure of atoms and the octet theory of valence. PNAS, 1919. f. Stable ion configurations - addition or subtraction of electrons from orbitals that yield electron configurations of the noble gases.- pseudo- noble gas configurations. i. Isoelectric species- atoms/ions that have similar electron configurations. g. Oxidation Number- net charge of the stable ion. i. Identifying oxidation states:

Periodic Variations of Oxidation Numbers 1.The maximum positive oxidation number for any representative element is equal to the group number, from +1 (Alkali metals) to +7 (Halogens). Noble Gases have an oxidation number of 0 2.Metallic elements usually exhibit only positive oxidation numbers due to low Zeff values 3.The most negative oxidation number for any representative element is equal to the group number minus eight. 4.Negative oxidation numbers are commonly limited to nonmetals. Semi-metals are negative only when these elements are combined with less electronegative elements. 5.Elements commonly exhibit positive oxidation numbers only when combined with more electronegative elements. 6.With the exception of Carbon, Nitrogen, Oxygen, and Mercury, each representative element that exhibits multiple oxidation numbers in its compounds commonly have oxidation numbers that are either all even or all odd.

II. Compounds a. Compound- a pure substance that is composed of two or more elements that are chemically bonded. b. Law of Definite Proportions- All samples of the same compound contain the same elements in the same proportion by mass. i. Coined by Joseph Proust. ii. John Dalton used this as a foundation to his Atomic Theory Calculating Percent Mass Percent Mass =Mass of the element x 100% Mass of the sample c. Law of Multiple Proportions. When two elements reacts to form more than one compound, the ratios of the element's masses. Can be reduced to a small whole number. i. Molecules- a group of atoms that are chemically bonded which act as an individual unit. 1. Typically formed from covalent bonds. d. Ionic Bond - chemical bond that occurs by a transfer of electrons and the resulting electrostatic attraction between the charged particles (ions). i. Due to larger differences in electronegativities (metal/nonmetal). 1. Electronegativity is the pull on electrons involved in a covalent bond. Derived by Linus Pauling. 2. Classifying Bond types based on electronegativity differences. a. Nonpolar covalent = 0.0 to 0.4 b. Moderately polar covalent = 0.4 to 1.0 c. Very polar covalent = 1.0 to 2.0 d. Ionic = greater than 2.0 ii. Ionic Compounds- Compounds formed from ionic bonds. 1. Characteristics: a. Brittle Crystalline Solids. b. High melting points- smaller ions produce higher melting points. c. Conduct electricity in the molten state (barring no decomposition).

d. Dissolve in high-polarity solvents (H2O) which results in conduction of electricity- electrolyte. e. Lattice Energy- the energy associated with bonded ions in a crystal. i. Lattice Energy is the energy needed to vaporize the ions in a compound. This is basically related to the amount of energy. Released when the ions come together forming the compound. Calculating Lattice Energy

Lattice Energy = k ( qc qa ) d q is the charge of the cation or anion d is the S of the ion radii

Print this off a table of Solubilities of Ionic Compounds e. Chemical formula- shows that kinds and numbers of atoms in a compound. i. Empirical formula- shows the simplest whole-number ration of atoms in a molecule. ii. Molecular formula- shows the actual ratio of elements in a molecule. iii. Formula Unit- lowest whole-number ratio of ions in an ionic compound. f. Percent Ionic Character (Pauling)- Based on D-electronegativities between cations and anions. i. Allen Factor (Leland C. Allen - Princeton)- used to predict Percent Ionic character. 1. High Allen Factors predict ionic compounds ( A.F. > 0.4). 2. Low Allen Factors predict covalent molecules (A.F. < 0.4) 3. Modified Allen Factor equation - Using Allen Factor to calculate percent ionic character: a. Percent Ion character = Allen Factor x 1.3 x 100% Calculating Allen Factor A.F. =  e.n. /  e.n. g. Writing & Naming Ionic Compounds i. Naming Cations 1. Cations formed from metals have the same name as the metal a. Single atom ions with or without multiple oxidative states are called monatomic ions 2. Cations that express multiple oxidation states (typically the transition metals) are named using: a. -ic/-ous- archaic scheme used to differentiate between an elements ions with only 2 oxidation states. b. Roman Numeral Nomenclature- the Roman numeral is used that corresponds to the oxidation state. c. Cations formed from nonmetal atoms have names that end in –ium. 3. Ions that are created from groups of covalently bonded atoms are called polyatomic ions a. Print off a list of Common Polyatomic Ions. ii. Naming Anions. 1. Monatomic anion name endings are changed to –ide. a. NOTE: There are some polyatomic anions that also end in -ide, so don't confuse them. 2. Polyatomic anions containing oxygen atoms are called oxyanions and have name endings of -ite or –ate. a. NOTE: -ate anions have 1 more oxygen atom than -ite anions. b. Some groups of anions need prefixes to differentiate between anions. (Ex. chlorates & chlorites). 3. Polyatomic anions that are made from adding an H+ to an oxyanion have the prefix hydrogen- or bi-. iii. Rules for Writing & Naming Ionic Compounds. 1. Cations are always written first. 2. The sum of the oxidation numbers must equal zero- Electrically neutral species. a. Use subscripts in order to get the correct proportion of atoms. b. The criss-cross method (oxidation #s to subscripts) works if you remember to reduce when necessary. 3. Name the compound using correct naming scheme for the cation and anion. a. Binary compounds- Composed of two different types of elements (metal/nonmetal). b. Ternary compounds- Composed of more than two types of elements (compounds containing polyatomic ions). c. NOTE: A polyatomic anion name ending does not change. III. Naming Acids. a. Acid- A compound which donates H+ ions in an aqueous solution. The hydrogen functions as the cation in the formula. i. Binary acid- hydrogen is bonded to a single non-metal and dissolved in water. 1. To name Binary Acids: a. use the prefix hydro- b. followed by the root of the non-metal c. end with the suffix -ic acid. d. Examples:

i. HCl(aq) -- hydrochloric acid

ii. HBr(aq) -- hydrobromic acid

iii. HF(aq) -- hydrofluoric acid

iv. HI(aq) -- hydroiodic acid

v. H2Se(aq) -- hydroselenic acid ii. Ternary acids- hydrogen is bonded to a polyatomic acid and dissolved in water. 1. To name Ternary Acids: a. DO NOT use the hydro- prefix b. use the root of the polyatomic ion c. change -ate endings to -ic acid or -ite ending to -ous acid d. Examples:

i. HNO3(aq) -- nitric acid

ii. HNO2(aq) -- nitrous acid

iii. HClO4(aq) -- perchloric acid

iv. HClO3(aq) -- chloric acid

v. HClO2(aq) -- chlorous acid

vi. HClO(aq) -- hypochlorous acid

Want to see another approach in the form of flowcharts? Download and print flowcharts for: Naming anions Naming cations Naming compounds with Acids Naming compounds