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Understanding Ionic Compounds: Key Concepts and Naming Conventions

This chapter covers essential knowledge regarding ionic compounds, including the proper writing of common ions (Fe²⁺, Fe³⁺) and polyatomic ions like sulfate (SO₄²⁻) and acetate (C₂H₃O₂⁻). It explains the formation of ionic bonds, stability through electron transfer, crystal lattice structure, and the characteristics of ionic compounds such as their melting points and electrical conductivity. The rules for naming ionic compounds, involving both monatomic and polyatomic ions, are detailed with examples. Homework exercises encourage practice and reinforcement of these concepts.

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Understanding Ionic Compounds: Key Concepts and Naming Conventions

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  1. Chapter 8 Ionic Compounds

  2. THINGS YOU MUST KNOW!!!!!!! • Pg 222 • Common Ions • Must be written correctly • Fe2+ Fe3+ (correct) • 2+, FE3+,Fe3+, Fe3+ (Incorrect) • Pg 224 Polyatomic Ions • Groups of covalently bonded atoms that have an overall charge and act like an ion. • SO42- Sulfate • What is C2H3O2-? • Acetate • What is NaC2H3O2? • Sodium acetate • On your quizzes you will be given the name. You must give the ion or formula.

  3. Ionic Bonds • Ion • Charged atom • Cations • Metals • Metals lose electrons and become positive • Anions • Nonmetals • Nonmetals gain electrons and become negative

  4. Why does bonding occur? • Every atom wants to be stable • Pseudonoble gas configuration • Pseudo means false • An atom that is not a noble gas but has the electron configuration of a noble gas. • This is achieved through bonding and makes the atoms involved stable.

  5. Rule: The sum of the charges for ionic compounds must equal 0. • Always written metal nonmetal • Examples: • Mg and Cl • You do: Rb and O • Do not swap charges to make subscripts.

  6. Ionic compounds form crystal lattices • 3-D repeating pattern of ions • Ions are bonded by lattice energy • Energy required to separate ions • Expressed as a negative value • The more negative the number the more energy required to separate the ions

  7. Characteristics of ionic compounds • High melting and boiling points • Hard • Strong bond • Nonconductors in solid state • Conductors in liquid state • Electrolytes • Ions in aqueous solution that conduct electricity • When ionic compounds come apart in a solution it is called dissociation

  8. Why do atoms gain and lose electrons? • Do ionic compounds form molecules? • No, because they form a crystal lattice. • Repeating pattern of ions based upon 3-D orientation of the way ions pack around one another. • The simplest ration of ions is called the formula unit. • Also why ionic compounds in a solid state are brittle, nonconductors, and hard. • Lattice energy is affected by 2 things • Size – The smaller the ion the greater the energy • Charge – The greater the charge the greater the energy

  9. Metallic bonds are similar to ionic bonds but they are not ionic bonds. • Metallic bonding is the reason metals have their properties. • Ductile, malleable, conduct electricity. • The outer shell of electrons are delocalized. • Not associated with one nucleus • Electron sea model • Used to illustrate metallic bonding

  10. Metals can form alloys • Mixture of metals • 2 types • Interstitial alloys • Smaller atom fills the interstasis • Substitutional alloys • Occurs when atoms are the same size

  11. Writing Ionic Compounds • Sum of charges must equal 0 • Metal Nonmetal • Ex. Strontium and Bromine • Sr2+ and Br- SrBr2 • Potassium and Oxygen • K+ and O2- K2O • You do: Aluminum and sulfur • You do: Calcium and arsenate

  12. Things You Must Know • 2NaCl • What does the 2 mean? • SrBr2 • What does the 2 mean?

  13. Naming Ionic Compounds • Metal and monatomic anion • Name of the metal and root of the anion with –ide ending. • Ex. Rb2S • Rubidium Sulfide • Metal w/ multiple charges and monatomic anion • Name of the metal with a roman numeral representing charge and root of the anion with –ide ending. • Ex. Copper and Bromine • Could be Cu+Br- or Cu2+Br- • Cu+Br- would be named copper (I) bromide • Cu2+Br- would be named copper (II) bromide

  14. Metal and polyatomic anion • Metal polyatomic anion • Ex. NaC2H3O2 • Sodium acetate • Metal w/ multiple charges and a polyatomic • Metal w/ charge and polyatomic • Ex. AuNO3 • Gold (I) nitrate • Fe(ClO)3 • Iron (III) hypochlorite

  15. Multiple Polyatomic Groups: Oxyanions • All contain oxygen • Two possibilities • -ate or –ite • -ate is used to show the polyatomic that contains the greater number of oxygen atoms • -ite is used to show the polyatomic that contains the leaser number of oxygen atoms. • Ex. NO3- is nitrate; NO2- is nitrite • Ex. SO42- is sulfate; SO32- is sulfite

  16. There are four possibilities with halogens that form a polyatomic anion. • Per- -ate -ate -ite Hypo- -ite • Ex. ClO4- is perchlorate ClO3- is chlorate ClO2- is chlorite ClO- is hypochlorite • You name: NaBrO

  17. Writing Formulas from Names • Aluminum selenide • Al3+ and Se2- Al2Se3 • Tin (IV) oxide • Sn4+ and O2-  SnO2 • Lead (II) hydrogen carbonate • Pb2+ and HCO3-  Pb(HCO3)2

  18. Homework • Pg. 237 74-84

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