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STAAR Chemistry Review Topic : Bonding

STAAR Chemistry Review Topic : Bonding. TEKS 7 – The student knows how atoms form ionic, covalent, and metallic bonds. 7A - E. INDEX CARD TIME! TITLE: Metals vs. Nonmetals FRONT: create a GUIDE on how to determine if an element is a metal or a nonmetal.

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STAAR Chemistry Review Topic : Bonding

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  1. STAAR Chemistry ReviewTopic: Bonding TEKS 7 – The student knows how atoms form ionic, covalent, and metallic bonds. 7A - E

  2. INDEX CARD TIME! TITLE: Metals vs. Nonmetals FRONT: create a GUIDE on how to determine if an element is a metal or a nonmetal. BACK: describe ionic, covalent and metallic bonds in terms of which types of elements (metals or nometals) are involved.

  3. Metals vs. Nonmetals on the P. T.

  4. 3 types of bonds • Ionic: between a METAL and a NONMETAL • Electrons are TRANSFERRED from the metal to the nonmetal • Covalent: between NONMETALS • Electrons are SHARED between the nonmetals • Metallic: between METALS • Electrons freely move around in a “SEA OF ELECTRONS”

  5. Student Expectation (SE) 7A – NAME ionic and covalent compounds using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules.

  6. INDEX CARD TIME! TITLE: Naming Ionic Compounds FRONT: create a GUIDE on how to name ionic compounds BACK: provide at least 3 EXAMPLES of naming ionic compounds Ionic = METAL + NONMETAL

  7. Mini-Review • Cations – Positively charged ions (metals) • Anions – Negatively charged ions (nonmetals) • Cations form by losing electrons • Anions form by gaining electrons • Charges for these ions are located at the top of the columns on the periodic table. (oxidation numbers)

  8. 0 Ionic: Naming 1. Name the METAL (cation) • group 1, group 2, silver, or zinc • simply use the element’s name. • Any other metal (transition metals): • Use the element’s name and the charge in ROMAN NUMERALSin parentheses. 2. Name the NONMETAL (anion) • single element: • use the element’s name with the –ide ending. • Polyatomic ion: • always use a polyatomic ion’s name unaltered. • NO GREEK PREFIXES! • Ex: Na2O = Sodium oxide • Fe2(SO4)3 = Iron (III) sulfate

  9. Examples • NaF: Sodium fluoride • KCl: Potassium chloride • Li2CO3: Lithium carbonate • Ca(NO3)2: Calcium nitrate • ZnSO4: Zinc sulfate • Cu(NO2)2: Copper (II) nitrite

  10. INDEX CARD TIME! TITLE: Naming Covalent Compounds FRONT: create a GUIDE on how to name covalent compounds BACK: provide at least 3 EXAMPLES of naming covalent compounds Covalent= 2 NONMETALS

  11. Mini-Review: Greek Prefixes • 1 mono- • 2 di- • 3 tri- • 4 tetra- • 5 penta- • 6 hexa- • 7 hepta- • 8 octa- • 9 nona- • 10 deca-

  12. 0 Covalent: Naming • Name the 1st element. -If there’s more than one, use a Greek Prefix. 2. Name the 2nd element. -Always use a Greek Prefix with the 2nd element. -change the ending to “–ide” • Ex: PCl3 = Phosphorus trichloride

  13. Examples • SO2: sulfur dioxide • SeF6: selenium hexafluoride • PCl5: phosphorus pentachloride • As2O5: diarsenicpentoxide • NO2: nitrogen dioxide • H2O: dihydrogen monoxide

  14. Student Expectation (SE) 7B – write the chemical formulas of common polyatomic ions, ionic compounds, and covalent compounds.

  15. INDEX CARD TIME! TITLE: Writing Formulas for Ionic Compounds FRONT: create a GUIDE on how to write formulas for ionic compounds BACK: provide at least 3 EXAMPLES. Ionic = METAL + NONMETAL

  16. 0 Ionic: Writing Formulas • Write the SYMBOLS for the ions. (with oxidation numbers) • Represent the oxidation number as a subscript on the opposite ion! • Use parentheses around multiple polyatomic ions. Ex: Magnesium Chloride = Mg+2 and Cl-1 -> MgCl2 Ex: Copper (II) sulfite = Cu+2 and (SO3)-2 Cu2(SO3)2 -> Cu(SO3)

  17. Step 1 – Write chemical symbols of the elements in a compound. • Example: Calcium oxide Ca O

  18. Step 2- Write the oxidation number for each element as a superscript. • Example: Calcium oxide Ca2+ O2-

  19. Step 3 – Criss-Cross superscripts to subscripts. • Example: Calcium oxide Ca2+ O2- Ca2 O2 = Ca2O2

  20. Step 4 – Reduce Subscripts if needed. • Example: Calcium oxide 2:2 ratio can be reduced to 1:1 CaO

  21. Transition Metal Cations • Because they have d-block electrons, transition metals often can form multiple ions. • It’s difficult to predict the charge on a transition metal cation. These must be supplied to you. • Ex. Naming: Fe2(SO4)3 = Iron (III) sulfate • Ex. Formula: Copper (II) sulfite = Cu+2 and (SO3)-2 • Cu2(SO3)2 -> Cu(SO3) These are on the Back of your Booklet!

  22. Examples • Aluminum bromide: AlBr3 • Lead (II) sulfide: PbS • Magnesium chlorate: Mg(ClO3)2 • Ammonium phosphate: (NH4)3PO4 • Iron (III) nitrite: Fe(NO2)3

  23. INDEX CARD TIME! TITLE: Writing Formulas for Covalent Compounds FRONT: create a GUIDE on how to write formulas for covalent compounds BACK: provide at least 3 EXAMPLES. Covalent = 2 NONMETALS

  24. Covalent: Writing Formulas • Write the SYMBOL for the 1st element. -Write a SUBSCRIPT telling how many there are according to the GREEK PREFIX. 2. Write the SYMBOL for the 2nd element. -Write a SUBSCRIPT telling how many there are according to the GREEK PREFIX. Ex: Dinitrogen trioxide = N2O3

  25. Examples • Diphosphoruspentoxide: P2O5 • Carbon tetrabromide: CBr4 • Silicon dioxide: SiO2 • Diarsenictrisulfide: As2S3 • Carbon monoxide: CO

  26. Student Expectation (SE) 7C – construct electron dot formulas to illustrate ionic and covalent bonds.

  27. INDEX CARD TIME! TITLE: IONIC electron dot structures FRONT: Provide electron dot structures for one metal and one nonmetal. BACK: show the resulting TRANSFER of electrons Ionic = METAL + NONMETAL

  28. Mini Review • The electron is physically TRANSFERRED from a metal to a nonmetal. Li F = LiF (Li+) (F-) Picture used courtesy of http://www.ider.herts.ac.uk/school/courseware/materials/bonding.html

  29. Electron Dot Structure • The valence electrons are indicated by dots placed around the element’s symbol. • This can be used to represent up to eight valence electrons for an atom. One dot is placed on each side before a second dot is placed on any side. “empty corner rule”. Octet Rule: atoms form bonds in order to get all 8 valence electrons.

  30. IONIC LEWIS DOT STRUCTURES • Not connected by bonds! But the transferred electrons (and the new charges) are shown . Al . . When ions pair, the total positive charge must balance out the total negative charge. This means that ionic compounds are electrically NEUTRAL.

  31. INDEX CARD TIME! TITLE: COVALENT electron dot structures FRONT: Provide electron dot structures for two different nonmetals. BACK: show the resulting SHARING of electrons with a Lewis Dot Structure Covalent = 2 NONMETALS

  32. Mini Review: Covalent Bond • The electron is SHARED between two nonmetals. They are true molecules. Cl Cl =Cl2 Picture used courtesy of http://www.ider.herts.ac.uk/school/courseware/materials/bonding.html

  33. COVALENT LEWIS DOT STRUCTURES • Single electrons (not in pairs) will form bonds with other single electrons on different atoms

  34. Student Expectation (SE) 7D –  describe the nature of metallic bonding and apply the theory to explain metallic properties such as thermal and electrical conductively, malleability, and ductility.

  35. INDEX CARD TIME! TITLE: Metallic Bonds FRONT: Describe the NATURE of a metallic bond BACK: List some PROPERTIES of metallic bonds Metallic = 2 or more METALS

  36. Nature of Metallic Bonds • Between METAL atoms. • The valence electrons of metal atoms can drift freely from one part of the metal to another- this is sometimes called a “sea of electrons” • Metallic bonds consist of the attraction between these free floating electrons and the positively charged metal ions (cations). This attraction is the “bond” that holds metals together.

  37. Properties of Metals • Lusterous- they are shiny! • High density- atoms are tightly packed. • Good conductors of electricity and heat. • Reason- electrons can flow freely.

  38. Properties of Metals • Ductile- they can be drawn into wires AND • Malleable- they can be hammered into shapes • Reason- cations can slide easily past each other because the sea of electrons insulates them and prevents strong repulsions.

  39. Student Expectation (SE) 7E – predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries using VSEPR.

  40. INDEX CARD TIME! TITLE: VSEPR FRONT: Write down the theory of VSEPR BACK: Describe how many bonding regions and how many lone pairs surround the central atom are needed for each VSEPR shape. VSEPR is for Covalent = 2 NONMETALS

  41. Mini-Review VSEPR: Valence-Shell Electron-Pair Repulsion • -Used for COVALENT compounds ONLY). • -Electron pairs form bonds between atoms • But they’re still negative, and they still try to repel each other. • -Bonds and Lone Pairs try to get as far apart from each other as possible. • -Shape of the molecule is determined by the number of bonding regions and lone pairs around the central atom.

  42. How to Determine VSEPR Shape • 1. Draw the Lewis Dot Structure. • 2. Count the # of bonding regions and # of lone pairs on central atom. • double/triple bonds = ONE bonding region • 3. Determine SHAPE

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