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CHEMICAL BONDING

CHEMICAL BONDING. Ionic Bonding Covalent Bonding Metallic Bonding. What is bonding ?. Joining of atoms to make compounds and molecules . Become more stable - lower potential energy - Fill valence shell (Octet Rule). Ionic Bonds. Covalent Bonds. Major Types of Bonding.

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CHEMICAL BONDING

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  1. CHEMICAL BONDING Ionic Bonding Covalent Bonding Metallic Bonding

  2. What is bonding? Joining of atoms to make compounds and molecules. • Become more stable - lower potential energy - Fill valence shell (Octet Rule) Ionic Bonds Covalent Bonds

  3. Major Types of Bonding METALLIC BONDING IONIC BONDING COVALENT BONDING

  4. Strength of Bonds 1. Ionic – Strongest form of bond with the highest difference in electronegativity • The greater the difference in electronegativity, the more ionic character the bond has • Difference in electronegativity 1.7 and above 2. Covalent – second strongest form of bond • The lower the difference in electronegativity, the more covalent character the bond has • Difference in electronegativity, 0-1.7 3. Metallic – lowest strength • Not based on electronegativity – simply held together by the electrostatic forces between delocalized electrons

  5. Substances with Metallic, Ionic, and Covalent Bonds

  6. Ionic Bonding • Ionic bonding involves the electrostatic force of attraction of a cation (+) and an anion (-). • Charges and amounts of atoms balance to become a neutral compound.

  7. Forms Crystals • Regular geometric patterns of the formula unit. (lattice structure) • Formula unit is simplest whole number unit.

  8. Ionic Bonding Characteristics • Formed between METALS AND NONMETALS • Solid at room temperature • Form Crystal lattice structures • Dissolve in water • Conduct electricity when melted or dissolved • High Melting points

  9. Forms of Chemical Bonds • There are 2 extreme forms of connecting or bonding atoms: • Ionic: completetransfer of electrons • Covalent: electrons equally shared Most bonds are somewhere in between.

  10. Electronegativity Attraction of the bonded electrons to one of the atoms. A B But… which atom?

  11. The one with the highest electronegativity! Electronegativity Trend Same as electron affinity and ionization energy.

  12. Bond Type = Difference in electronegativities Greater than 1.7 = Ionic Bond 1.7 - .3= Polar Covalent Bond .3 – 0= Non Polar Covalent Bond SUBTRACT!

  13. Practice • What type of bond exists between sodium and oxygen? IONIC BOND 3.5 – 0.9 = 2.6 Greater than 1.7 = Ionic 1.7 - .3 = Polar Covalent .3 and below = NonPolar Covalent

  14. Bond type is on sliding scale of ionic & covalent bond extremes known as % ionic character. Pure Covalent Bond 1.7 (50%) Electroneg. Difference Ionic Bond

  15. Electronegativity • Due to it being a sliding scale the greater the electronegativity difference, the more ionic the bond. • Label the bond types formed below • Na – I • N – O • Al - Cl

  16. ELECTRON DOTS & IONS For atoms to enter into chemical reactions and become more stable, electrons are lost or gained and IONS are formed. Atoms will gain or lose electrons from the principal quantum level in order to achieve a noble gas configuration. With the exception of helium, 8 electrons in the outer energy level make an element stable and unreactive. This is the OCTET RULE.

  17. Electron Dot Structures • A visual way to show the valence electrons • Follow the orbital filling rules. • Example: Sulfur S electrons Electron Configuration: 1s22s22p63s23p4 S pz electrons px electrons py electrons

  18. You Try! Write structures for H, Mg, N, Cl, Ne H N Mg Cl Ne

  19. IONS Atoms that have gained or lost electrons and are no longer neutral. Atoms gain or lose electrons to become more stable (lower potential energy). They become like a noble gas configuration. Follows Octet Rule!

  20. IONS Electron Configuration: 1s22s22p63s1 Na Examples: Sodium Choice… Lose 1 electron and drop to full 2nd energy level or…. Gain 7 electrons to get a full 3rd energy level? But: (+)11 protons (-) 10 electrons (+) 1 +1 Na New electron configuration: 1s22s22p6 (Like Neon)

  21. IONS Electron Configuration: 1s22s22p63s23p4 S Examples: Sulfur Choice… Lose 6 electrons and drop to full 2nd energy level or…. Gain 2 electrons to get a full 3rd energy level? But: (+)16 protons (-) 18 electrons (-) 2 -2 S New electron configuration: 1s22s22p63s23p6(Like Argon)

  22. OXIDATION • Process of an element becoming more POSITIVE. • Could occur from neutral to +1 or from +1 to +2. • How does element oxidize? • Loss of electrons! Example: Na0Na1++ 1e- The ionic charge of Na (0 or 1+) can also be called its oxidation number.

  23. REDUCTION • Process of an element becoming more NEGATIVE. • Could occur from neutral to -1 or from -1 to -2. • How does element reduce? • Gain of electrons! Example: O0 + 2e-O2- What’s the oxidation number for O?

  24. Practice • Write the oxidation or reduction for the following and label as oxidation or reduction: • Aluminum • Iodine • Nitrogen • Beryllium

  25. Lewis Dot Diagrams for Ionic Bonding Mg2+ + 2 Br  MgBr2

  26. Parts of an Equation REACTANTSOXIDATION NUMBER PRODUCTSCOEFFICIENT SUBSCRIPT Mg2+ + 2 Br  MgBr2 *Cation is always written first! *Atoms combine in simplest whole number ratios to make their charges zero.

  27. You Try! Write the Lewis dot equation showing the ionic bonding between oxygenandsodium. 2Na+ + Br  Na2O

  28. PRACTICE Write the oxidation number (ionic symbols) of the following elements: Ba, I, Br, Be Write oxidation / reduction equations for aluminum and oxygen forming ions. Write the Lewis dot equation combining aluminum and oxygen.

  29. Naming Ionic Compounds Terms Monatomic ion Na1+or O2- Polyatomic ion NO31-NH4+ oxyanion– polyatomic w/ Oxygen Binary Ionic Compound Na2O Ternary Ionic Compound NaNO3 IUPAC: International Union of Pure and Applied Chemistry

  30. Naming Binary Compounds (2 different elements only) • Ex. FeCl3 – Iron (III) Chloride 1. Name the cation (positive ion) in the formula with no change. For the above example: Iron. 2. Determine if the cation is a transition metal or underneath the stair-step line (except cadmium, silver, aluminum and zinc). If it isn’t skip to step 5. For the above example: Iron IS a transition metal so go on to step 3. 3. Use the charge of the anion to determine the charge of the Iron since it is a transition metal. Since chlorine is a 1- charge and there are 3 of them, Iron has to be a 3+ in this case to balance it out. 4. Place the charge, in parenthesis, after the name of the cation. 5. Name the anion with an –ide ending.

  31. Ternary compounds (Containing a polyatomic ion) • Ex. Cr3(PO4)2 – Chromium (II) Phosphate 1. Name the cation in the formula with no change. For the above example: Chromium. 2. Determine if the cation is a transition metal or underneath the stair-step line (except cadmium, silver, aluminum and zinc). If it isn’t skip to step 5, if it is go to step 3. For the above example: Chromium IS a transition metal. 3. Use the charge of the polyatomic anion to determine the charge of the Chromium since it is a transition metal. Since phosphate is a 3- charge and there are 2 of them, Chromium has to be a 2+ in this case to balance it out. 4. Place the charge, in parenthesis, after the name of the cation. 5. Name the polyatomic ion with not change.

  32. Examples • Write the name of the following compounds: • Na2O ___________________________ • Cu3N2 ___________________________ • PbO ___________________________ • SnO2 ___________________________ • Na2CO3 ___________________________ • Mg(ClO2)2 ___________________________

  33. Practice • Naming compounds on whiteboards

  34. Binary Formulas • Example – Magnesium Chloride 1. Write down the symbol and charge of each ion, always placing the positive one first. Mg 2+Cl1- 2. Figure out how many of each ion you need to balance their charges and make them equal zero when put together as a compound. (in this case 1 magnesium, 2 chlorines) 3. Use subscripts to indicate how many of each ion you have in the formula. MgCl2

  35. Ternary Formulas • Example – Calcium Phospate 1. Write down the symbol and charge of each ion, always placing the positive one first. Ca2+ PO43- 2. Figure out how many of each ion you need to balance their charges and make them equal zero when put together as a compound (in this case 3 calciums and 2 phosphates) 3. Use subscripts to indicate how many of each ion you have in the formula, placing the polyatomic ion (PO4) in parenthesis (when more than one) with the subscript outside of the parenthesis. Ca3(PO4)2

  36. Examples • Write the formula for the following compounds: sodium oxide _______________ tin (IV) chloride ________________ nickel (II) phosphide _____________ ammonium hydroxide _____________ magnesium acetate________________

  37. Whiteboard Practice

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