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Review: What do we already know?

Review: What do we already know?. Protons determine the IDENTINTY of the element Valance Electrons determine the CHEMICAL properties of an element. Valance electrons are the electrons in the OUTER ENERGY level

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Review: What do we already know?

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  1. Review: What do we already know? • Protons determine the IDENTINTY of the element • Valance Electrons determine the CHEMICAL properties of an element. • Valance electrons are the electrons in the OUTER ENERGY level • For representative elements look at the group number to determine the number of valance electrons • Elements lose or gain electrons to achieve a full outer energy level (full OCTET) • Metals form CATIONS by losing electrons • Nonmetals form ANIONS by gaining electrons

  2. Review:The Octet Rule • Octet rule stated that in forming compounds atoms tend to achieve the electron configuration of a noble gas. • An octet is a set of eight electrons • Atoms of metals tend to lose their valence electrons leaving a complete octet in the next-lowest energy level. • Atoms of some nonmetals tend to gain electrons or to share electrons with another nonmetals to achieve a complete octet. • What is a compound?

  3. bonding Chemistry Chapter 5 & 6 I. Introduction to bonding

  4. Bonding:A. vocabulary • Chemical Bond • attractive force between atoms or ions that binds them together as a unit • bonds form in order to… • decrease potential energy (PE) • increase stability

  5. B. Types of Bonds COVALENT IONIC e- are transferred from metal to nonmetal e- are shared between two nonmetals Bond Formation Type of Structure true molecules crystal lattice Physical State liquid or gas solid Melting Point low high Solubility in Water yes usually not yes (solution or liquid) Electrical Conductivity no Other Properties odorous

  6. B. Types of Bonds METALLIC e- are delocalized among metal atoms Bond Formation Type of Structure “electron sea” Physical State solid Melting Point very high Solubility in Water no yes (any form) Electrical Conductivity malleable, ductile, lustrous Other Properties

  7. Most bonds are a blend of ionic and covalent characteristics. Difference in electronegativity determines bond type. C. Bond Polarity

  8. Electronegativity Attraction an atom has for a shared pair of electrons. higher e-neg atom  - lower e-neg atom + C. Bond Polarity

  9. Electronegativity Trend (p. 151) Increases up and to the right. C. Bond Polarity

  10. Nonpolar Covalent Bond e- are shared equally symmetrical e- density usually identical atoms C. Bond Polarity

  11. Polar Covalent Bond e- are shared unequally asymmetrical e- density results in partial charges (dipole) - + C. Bond Polarity

  12. C. Bond Polarity • Nonpolar • Polar • Ionic View Bonding Animations.

  13. C. Bond Polarity Examples: • Cl2 • HCl • NaCl 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9 Polar 3.0-0.9=2.1 Ionic

  14. II. Ionic Bonding and Naming Chapter 7 and 9

  15. SC1 Students will analyze the nature of matter and its classifications. • SC1.b. Identify substances based on chemical and physical properties. • SC1.c. Predict formulas for stable ionic compounds (binary and tertiary) based on balance of charges. • SC1.d. Use IUPAC nomenclature for both chemical names and formulas: • SC1.d.1 Ionic compounds (Binary and tertiary) • SC1.d.3 Acidic compounds (Binary and tertiary) • SC3.e. Compare and contrast types of chemical bonds (i.e. ionic, covalent). • SC3.b. Use the orbital configuration of neutral atoms to explain its effect on the atom’s chemical properties.

  16. C. Ionic Nomenclature 1+ 0 Common Ion Charges 2+ 3+ NA 3- 2- 1-

  17. A. Vocabulary ION 2 or more atoms 1 atom Monatomic Ion Polyatomic Ion Na+ NO3-

  18. Lewis Structure • Electron dot structures (Lewis dot structure) are diagram that show the valence electrons at dots. • Each valance electron is represented with a dot • Put one dot on each side of the symbol before putting two on one side. • Examples Carbon Calcium Chlorine Argon 4 valance e- 2 valance e- 7 valance e- 8 valance e- C Ca Cl Ar

  19. Ionic – show transfer of e- B. Lewis Structures

  20. Formation of Ionic Compounds • Ionic compounds are compounds composed of cations and anions. • Although they are compounds of ions, ionic compounds are electrically neutral. • Ionic bonds are the electrostatic forces that hold ions together in ionic compounds. They occur due to the transfer of electrons • Chemical formula shows the kinds and numbers of atoms in the smallest representative unit of a substance. • Formula unit is the lowest whole-number ratio of ions in an ionic compound

  21. A. Vocabulary COMPOUND more than 2 elements 2 elements Binary Compound Ternary Compound NaCl NaNO3

  22. Writing Ionic Formulas Write each ion, cation first. Don’t show charges in the final formula. Overall charge must equal zero. If charges cancel, just write symbols. If not, use subscripts to balance charges. Use parentheses to show more than one polyatomic ion. Stock System - Roman numerals indicate the ion’s charge. C. Ionic Nomenclature

  23. Formula Writing • When writing formulas the MOST METALIC element is written first • Ionic bonds occur between METALS and NONMETALS so the metal is ALWAYS written FIRST. • Determine the ion that the elements will form • Balance charges • Can switch charges and reduce if necessary • Or can use the following equation: # of metal x charge of metal + # of nonmetal x charge of nonmetal = 0 atoms atoms

  24. Formula Writing • Practice 1: Oxygen and Sodium • Sodium is metal so it MUST be written first • Na forms +1 ion and O forms -2 ion • Na+1 O-2 switching charges gives Na2O (# metal) +1 +(#nonmetal) -2 = 0 solve (2) +1 +(1) -2 = 0 gives Na2O NOTE: The subscript of 1 is NOT written • Practice 2: Nitrogen and Aluminum • Aluminum is metal so it MUST be written first • Al forms +3 ion and N forms -3 ion • Al+3 N-3 switching charges gives AlN (must reduce) (# metal)+3 +(#nonmetal) -3 = 0 solve (1) +3 +(1) -3 = 0 gives AlN

  25. Formula Writing • Practice 3: Calcium and Carbon • Calcium is metal so it MUST be written first • Caforms +2 ion and C forms -4 ion • Ca+2 C-4 switching charges gives Ca2C (# metal)+2 +(#nonmetal) -4 = 0 solve (2)+2 +(1) -4 = 0 gives Ca2C • Practice 4: Barium and Phosphate (PO4-3) • Barium is metal so it MUST be written first • Ba forms +2 ion and PO4 is a -3 ion • Al+2PO4-3 switching charges gives Al3(PO4)2 (# metal)+2 +(#nonmetal) -3 = 0 solve (3) +2 +(2) -3 = 0 gives Al3(PO4)2 MUST USE parenthesis to show having 2 phosphate molecules.

  26. C. Ionic Nomenclature Ionic Names • Write the names of both ions, cation first. • Change ending of monatomic ions to -ide. • Polyatomic ions have special names. • Stock System - Use Romannumerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.

  27. C. Ionic Nomenclature • Consider the following: • Does it contain a polyatomic ion? • -ide, 2 elements  no • -ate, -ite, 3+ elements  yes • Does it contain a Roman numeral? • Check the table for metals not in Groups 1 or 2. • No prefixes!

  28. C. Ionic Nomenclature 1+ 0 Common Ion Charges 2+ 3+ NA 3- 2- 1-

  29. C. Ionic Nomenclature • potassium chloride • magnesium nitrate • copper(II) chloride  KCl • K+ Cl- • Mg2+ NO3-  Mg(NO3)2  CuCl2 • Cu2+ Cl-

  30. C. Ionic Nomenclature • NaBr • Na2CO3 • FeCl3 • sodium bromide • sodium carbonate • iron(III) chloride

  31. Ionic Naming – Type 1 metals Type 1 metals are metals that form only 1 oxidation state. They are found in groups 1, 2, & 13 in addition to Zn+2, Cd+2, and Ag+1 • Determine the Cation and the Anion • If the cation is from a representative element write its name • The anion is a: • Polyatomic ion write it’s special name • Single nonmetal element write its root name followed by “ide” • write both parts of name side by side

  32. Ionic Naming – Type 1 • Example 1: Ca3N2 • Calcium is the cation, nitrogen is anion • Calcium stays calcium • Nitrogen is NOT a polyatomic so it becomes Nitride • Ca3N2 is called calcium nitride • Example 1: Ca3(PO4)2 • Calcium is the cation, phosphate is anion • Calcium stays calcium • Phosphate is a polyatomic so it’s name is phosphate • Ca3(PO4)2is called calcium phosphate

  33. Ionic Naming – Type 1 • Practice • Potassium chloride • Aluminum chloride • Calcium carbide • Indium nitride • Rubidium phosphite • Aluminum hydroxide • Indium sulfite • ** Ammonium bromide • Practice • KCl • AlCl3 • Ca2C • InN • Rb3PO3 • Al(OH)3 • In2(SO3)3 • (NH4)3Br

  34. Ionic Naming – Type 2 metals • Most transition metals have the ability to borrow electrons from other orbitals and can form ions with different charges. • Metals in group 14 also have multiple oxidation states. +2 or +4 • Example: Iron can from a +3 or +4 cation, copper can from a +2 or +1 ion • Not ALL transition metal do this but MOST do so when we name the compound we have to state the charge of the metal ion • EXCEPTIONS: three transition metals that you MUST memorize the following: Zn+2, Cd+2, Ag+1 as they do NOT need roman numerals

  35. Ionic Naming – Type 2 Metals • Determine the Cation and the Anion • If the cation is from a transition element write its name followed by a roman numeral to show the charge of the metal ion. • The anion is a: • Polyatomic ion write it’s special name • Single nonmetal element write its root name followed by “ide” • write both parts of name side by side

  36. Ionic Naming – Type 2 • Common Roman numerals you MUST KNOW • I 6. VI ** • II 7. VII • III 8. VIII • IV ** 9. IX • V 10. X ** commonly confused by students

  37. Ionic Naming – Type 2 • Example 1: FeO • Fe is cation and O is anion • since Oxygen has a -2 charge Fe MUST have a +2 so it is Iron (II) • O is NOT a polyatomic so it becomes Oxide • FeO is Iron (II) Oxide • Example 2: Fe2O3 • Fe is cation and O is anion • since Oxygen has a -2 charge Fe MUST have a +3 so it is Iron (III) • O is NOT a polyatomic so it becomes Oxide • FeO is Iron (III) Oxide

  38. Ionic Naming – Type 2 • Example 3: FePO4 • Fe is cation and PO4 is anion • since PO4 has a -3 charge Fe MUST have a +3 so it is Iron (III) • PO4 is a polyatomic so it is phosphate • FeO is Iron (III) Phosphate • Example 4: Ag2O • Ag is cation and O is anion • Ag is an exception and only forms a +1 ion so is Silver • O is NOT a polyatomic so it becomes Oxide • Ag2Ois Sliver Oxide Remember Zn+2, Cd+2, Ag+1do NOT need roman numerals

  39. C. Ionic Nomenclature Ionic Formulas • Write each ion, cation first. Don’t show charges in the final formula. • Overall charge must equal zero. • If charges cancel, just write symbols. • If not, use subscripts to balance charges. • Use parentheses to show more than one polyatomic ion. • Stock System - Roman numerals indicate the ion’s charge.

  40. Ionic Names to formula • Use the name to determine the ions of the elements (or polyatomic) in compound • Write the ions for each element • Balance charges using either method(reduce if necessary) • Chemical formulas for COMPOUNDS do NOT have charges!! • The number of atoms MUST be shown as a subscript. • REMEMBER the size and the shape of the letters matter when writing chemical formulas: COS and CoS are two different things

  41. Ionic Names to Formulas • Examples • Zinc Oxide • Zn+1 O-2 • Zn2O • Cobalt (II) Oxide • Co+2 O-2 • CoO • Manganese (IV) Sulfate • Mn+4 S-2 • MnS2 • Examples • Strontium Sulfide • Sr+2 S-2 • SrS • Magnesium Cyanide • Mg +2 CN-1 • Mg(CN)2 • Potassium Phosphide • K+1 P+3 • K3P

  42. Formation of Ionic Compounds • Metals lose their electrons to nonmetals • The opposite charges attract and form an ionic bond Na + Cl →Na + Cl →Na+1 + Cl-1 → NaCl name is: sodium chloride Mg + S →Mg + S →Mg+2 + S-2 → MgS name is: Magnesium Sulfide

  43. Properties of Ionic Compounds • Ionic compounds form by the TRANSFER of electrons • Most ionic compounds are crystalline solids at room temperature. • Ions in the crystals are arranged in repeating three-dimensional patterns. • The large attractive forces result in very stable structures • Ionic compounds generally have high melting points. • Ionic compounds can conduct an electric current when melted or dissolved in water • The ion movement allows electricity to flow between electrodes

  44. Metallic Bonds and Metallic Properties • The valence electrons of metal atoms can be modeled as a sea of electrons. • Metallic bonds consist of the attraction of the free-floating valence electrons for the positively charged metal ions. • the sea-of-electrons models explains many physical properties of metals • metals are good conductors of electric current because electrons can flow freely in them • metals are malleable (can be hammered or forced into shapes.) • metals are ductile (can be drawn into wires)

  45. Alloys • metal atoms are arranged in very compact and orderly patterns • alloys are mixtures composed of two or more elements, at least of one which is metal • alloys are important because their properties are superior to those of their components elements. • Bronze alloy is made of copper and iron • Steel alloys are made of iron and carbon with additional elements.

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