CHEM 98 Ionic Compounds Lecture Notes PDF

CHEM 98 Ionic Compounds  Outline Reference Fundamentals of General, Organic, and Biological Chemistry 8th Edition McMurry, Ballantine, Hoeger, Peterson Chapter 3 3.1 Ions 3.2 Ions and the Octet Rule 3.3 Ions of Some Common Elements 3.4 Periodic Properties and Ion Formatio...

CHEM 98 Ionic Compounds  Outline Reference Fundamentals of General, Organic, and Biological Chemistry 8th Edition McMurry, Ballantine, Hoeger, Peterson Chapter 3 3.1 Ions 3.2 Ions and the Octet Rule 3.3 Ions of Some Common Elements 3.4 Periodic Properties and Ion Formation 3.5 Naming Monoatomic Ions 3.6 Polyatomic Ions 3.7 Ionic Bonds 3.8 Formulas of Ionic Compounds 3.9 Naming Ionic Compounds 3.10 Some Properties of Ionic Compounds 3.11 H+ and OH– Ions: An Introduction to Acids and Bases  3.1 Ions Ion – Formed when a neutral atom gains or loses electrons. – The loss of one or more electrons gives a positively charged ion called a cation. – The gain of one or more electrons gives a negatively charged ion called an anion.  3.1 Ions Alkali metals (group 1A) form compounds with halogens (group 7A). Properties of these compounds include the following: – High melting points – Stable, white, crystalline solids – Soluble in water – Conduct electricity when dissolved in water  3.1 Ions Solution of sodium chloride in water Conducts electricity Allows the bulb to light Electricity can only flow through a medium containing charged particles (ions) that are free to move.  3.1 Ions Sodium and other alkali metal atoms – Single electron in their valence shell – Electron configuration of ns1. – By losing this electron, an alkali metal is converted to a positively charged cation with a stable noble gas configuration. Same electron configuration as neon  3.1 Ions Chlorine and other halogen atoms – 7 valence electrons and have a ns2np5 electron configuration. – By gaining an electron, a halogen is converted to a negatively charged anion with a stable noble gas configuration. Same electron configuration as argon 2.8 Electron Configurations and the Periodic Table  3.1 Ions The symbol for a cation is written by adding the positive charge as a superscript to the symbol for the element. – The sodium cation would be written as Na+ The symbol for an anion is written by adding the negative charge as a superscript. – The chlorine anion would be written as Cl– If the charge is greater than 1, the number is used, as in Ca2+ & N3–  3.2 Ions and the Octet Rule Octet rule: Main group elements tend to undergo reactions that leave them with eight valence electrons. – Full valence shell  3.2 Ions and the Octet Rule When sodium (or any other alkali metal) reacts with chlorine (or any other halogen): The metal transfers an electron from its valence shell to the valence shell of the halogen. Both atoms then achieve ‘full’ outer valence shells  3.2 Ions and the Octet Rule Main group metals – Lose electrons to form cations – Attain an electron configuration like that of the noble gas just before them in the periodic table. Main group nonmetals – Gain electrons to form anions – Attain electron configuration like that of the noble gas just after them in the periodic table.  Worked Example 3.1 Write the electron configuration of magnesium (Z=12) Show how many electrons a magnesium atom must lose to form an ion with a filled shell (eight electrons) Write the configuration of the ion Explain the reason for the ion’s charge, and write the ion’s symbol  Worked Example 3.2 How many electrons must a nitrogen atom, Z=7, gain to attain a noble gas configuration? Write the electron dot and ion symbols for the ion formed. 3.3 Ions of Some Common Elements  3.3 Ions of Some Common Elements Group 1A: M → M+ + e– Group 2A: M → M2+ + 2e– Group 3A: Al → Al3+ + 3e–; no other common ions Group 4A, 5A: No common ions Group 6A: X + 2e– → X2– Group 7A: X + e– → X– Group 4A, 5A – Transition metals form cations – But they can lose one or more d electrons in addition to losing valence s electrons. – The octet rule is not followed.  3.3 Ions of Some Common Elements Metals form cations by losing one or more electrons. – Group 1A metals form +1 ions – Group 2A metals form +2 ions – Both achieve a noble gas configuration – Transition metals can form cations of more than one charge by losing a combination of valence-shell s electrons and inner-shell d electrons.  3.3 Ions of Some Common Elements Reactive nonmetals form anions by gaining one or more electrons to achieve a noble gas configuration. – Group 6A nonmetals oxygen and sulfur form the anions O2– and S2–. – Group 7A elements (the halogens) form –1 ions. Group 8A elements (the noble gases) are unreactive.  3.3 Ions of Some Common Elements Ionic charges of main group elements can be predicted using the group number and the octet rule. – For 1A and 2A metals Cation (+) charge = group number. – For nonmetals in groups 5A, 6A, and 7A Anion charge = 8 – (group number).  Worked Example 3.3 Which of the following ions is likely to form? (a) S3– (b) Si2+ (c) Sr2+  3.4 Periodic Properties and Ion Formation Metals Left side of the periodic table tend to lose electrons Nonmetals Right side of the periodic table tend to gain electrons. How does this trend change as you move across the periodic table?  3.4 Periodic Properties and Ion Formation Ionization energy – Energy required to remove one electron from a single atom in the gaseous state. – Small values indicate ease of losing electrons to form cations. Electron affinity – Energy released on adding an electron to a single atom in the gaseous state. – Halogens have the largest values and gain electrons most easily.  3.4 Periodic Properties and Ion Formation Halogens gain electrons most easily. Alkali metals lose electrons most easily. Noble gases neither lose nor gain electrons.  Worked Example 3.5 Which element is likely to lose an electron more easily, Mg or S?  3.5 Naming Monoatomic Ions Naming Cations – Main group metal cations are named by: – Identifying the metal, followed by the word ion. K+ Mg2+ Al3+ Potassium ion Magnesium ion Aluminum ion  3.5 Naming Monoatomic Ions Transition metals can form more than one type of cation. Two naming systems are used. – Old: The ion with the smaller charge is given the ending -ous, and the ion with the larger charge is given the ending -ic. – New: The charge on the ion is given as a roman numeral in parentheses right after the metal name. Cr2+ Cr3+ Old name: Chromous ion Chromic ion New name: Chromium (II) ion Chromium (III) ion 3.5 Naming Monoatomic Ions  3.5 Naming Monoatomic Ions Naming Anions – Replace the ending of the element name with -ide, followed by the word ion.  3.6 Polyatomic Ions Polyatomic ion: Ion composed of more than one atom  3.6 Polyatomic Ions Atoms in polyatomic ions held together by covalent bonds. – More on covalent bonds next chapter A polyatomic ion is charged because: – Contains a total number of electrons that is different from the total number of protons in the combined atoms. – # of electrons ≠ # of protons  3.7 Ionic Bonds Sodium chloride crystal Arrangement of Na+ and Cl– ions Crystal is held together by ionic bonds.  3.7 Ionic Bonds Ion-transfer reactions of metals and nonmetals form products unlike either element. (a) Chlorine – toxic green gas, sodium – reactive metal, sodium chloride – ‘harmless’ white solid (b) Sodium metal burning when immersed in chlorine gas, yielding white sodium chloride ‘smoke’  3.7 Ionic Bonds Ionic compounds Opposite electrical charges attract each other – Positive ion and negative ion are held together by an ionic bond. Crystals are ionic solids – Many ions attracted by ionic bonds to their nearest neighbors.  3.8 Formulas of Ionic Compounds All chemical compounds are neutral Determining Formulas of Ionic Compounds 1. Identify the ions 2. Decide how many ions of each type give a total charge of zero. Chemical formula of an ionic compound tells the ratio of anions and cations.  3.8 Formulas of Ionic Compounds If ions have the same charge number – One of each is needed. K+ + F– → KF If ions have different charge numbers – Unequal numbers of anions and cations must combine to have a net charge of zero. 2 K+ + O2– → K2O Ca2+ + 2 Cl– → CaCl2  3.8 Formulas of Ionic Compounds When the two ions have different charges: – Number of one ion is equal to the charge on the other ion (and vice versa) – ‘Crossover’  3.8 Formulas of Ionic Compounds Formula unit: – Formula that identifies the smallest neutral unit of an ionic compound – NaCl Formula unit is one Na+ ion and one Cl– ion. – CaF2, Formula unit is one Ca2+ ion and two F– ions. 3.8 Formulas of Ionic Compounds  3.8 Formulas of Ionic Compounds Rules for Writing Formula Once the numbers and kinds of ions in a compound are known – Formula is written using the following rules: List the cation first and the anion second. Do not write the charges of the ions. Use parentheses around a polyatomic ion formula if it has a subscript.  Worked Example 3.6 Write the formula for the compound formed by calcium ions and nitrate ions.  3.9 Naming Ionic Compounds Naming Ionic Compounds Cite cation first: – Sodium Anion second – Chloride Space between the words – Sodium Chloride  3.9 Naming Ionic Compounds There are two kinds of ionic compounds: – Type I Contain cations of main group elements. – Type II Contain metals that can exhibit more than one charge. – Require different naming conventions.  3.9 Naming Ionic Compounds Type I – Contain cations of main group elements. – The charges on these cations do not vary. – Do not specify the charge on the cation. NaCl is sodium chloride. MgCO3 is magnesium carbonate.  3.9 Naming Ionic Compounds Type II – Contain metals that can exhibit more than one charge. – Specify the charge on the cation Old System: (-ous, -ic) New System: (roman numerals) FeCl2 is iron(II) chloride or ferrous chloride. FeCl3 is iron(III) chloride or ferric chloride.  3.9 Naming Ionic Compounds Type II ionic compounds continued: – Do not name: – FeCl2 iron dichloride – FeCl3 iron trichloride. – Once the charge on the metal (cation) is known, the number of anions needed to yield a neutral compound is also known. – Iron (II) chloride (1- charge on a Cl ion, so 2 Cl- ions) – Iron (III) chloride (1- charge on a Cl ion, so 3 Cl- ions) – Charges do not need to be included as part of the compound name.  Worked Example 3.7 Sodium and calcium both form a wide variety of ionic compounds. Write formulas for the following compounds. (a) Sodium bromide and calcium bromide (b) Sodium sulfide and calcium sulfide (c) Sodium sulfate and calcium sulfate (d) Sodium phosphate and calcium phosphate  Worked Example 3.8 Name the following compounds using roman numerals to indicate the charges on the cations where necessary. (a) KF (b) MgCl2 (c) AuCl3 (d) Fe2O3  3.10 Some Properties of Ionic Compounds The melting point of sodium chloride is 801 °C.  3.10 Some Properties of Ionic Compounds Ions in each compound settle into a pattern that efficiently fills space and maximizes ionic bonding. Ions in an ionic solid are held rigidly in place by attraction to their neighbors.  3.10 Some Properties of Ionic Compounds Once an ionic solid is dissolved in water – Ions can move freely – Accounts for the electrical conductivity of these compounds in solution.  3.10 Some Properties of Ionic Compounds Ionic compounds – Very high melting and boiling points. – Sodium chloride melts at 801 °C and boils at 1413 °C. Ionic compounds dissolve in water – If the attraction between water and the ions overcomes the attraction of the ions for one another – Not all ionic compounds are water soluble.  3.11 H+ and OH– Ions: An Introduction to Acids and Bases Two of the most important ions – Hydrogen cation (H+) – Hydroxide anion (OH–). They are fundamental to the concepts of acids and bases.  3.11 H+ and OH– Ions: An Introduction to Acids and Bases Acid – Substance that provides H+ ions in water – A hydrogen cation is simply a proton. When an acid dissolves in water – Proton attaches to a molecule of water to form a hydronium ion. H+ + H2O → H3O+ Chemists use the terms hydrogen and hydronium ions interchangeably.  3.11 H+ and OH– Ions: An Introduction to Acids and Bases Different acids can provide different numbers of H+ ions per acid molecule. – Hydrochloric acid, HCl Provides one H+ ion per acid molecule. – Sulfuric acid, H2SO4 Can provide two H+ ions per acid molecule. – Phosphoric acid, H3PO4 Can provide three H+ ions per acid molecule. © 2017 Pearson Education, Inc.  3.11 H+ and OH– Ions: An Introduction to Acids and Bases Base – Substance that provides OH– ions in water A hydroxide anion – Polyatomic ion – Oxygen atom covalently bonded to hydrogen atom.  3.11 H+ and OH– Ions: An Introduction to Acids and Bases Bases – Sodium hydroxide (NaOH) – Potassium hydroxide (KOH) When a base dissolves in water – OH– anions go into solution along with the metal cation.  3.11 H+ and OH– Ions: An Introduction to Acids and Bases Different bases can provide different numbers of OH– ions per formula unit. – Sodium hydroxide, NaOH Provides one OH– ion per formula unit. – Barium hydroxide, Ba(OH)2 Can provide two OH– ions per formula unit.  3.11 H+ and OH– Ions: An Introduction to Acids and Bases