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Chemical Bonds

Chemical Bonds. __________________ = atoms tend to gain, lose or share electrons so as to have 8 electrons. Gain 4 electrons. C would like to N would like to O would like to. Gain 3 electrons. Gain 2 electrons. Electron Dot diagrams are….

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Chemical Bonds

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  1. Chemical Bonds

  2. __________________ = atoms tend to gain, lose or share electrons so as to have 8 electrons Gain 4 electrons • C would like to • N would like to • O would like to Gain 3 electrons Gain 2 electrons

  3. Electron Dot diagrams are… • A way of showing & keeping track of valence electrons. • How to write them? • Write the symbol - it represents the nucleus and inner (core) electrons • Put one dot for each valence electron (8 maximum) • They don’t pair up until they have to (Hund’s rule) X

  4. The Electron Dot diagram for Nitrogen • Nitrogen has 5 valence electrons to show. • First we write the symbol. N • Then add 1 electron at a time to each side. • Now they are forced to pair up. • We have now written the electron dot • diagram for Nitrogen.

  5. Electron Dot Structures Symbols of atoms with dots to represent the valence-shell electrons

  6. Learning Check  A. X would be the electron dot formula for 1) Na 2) K 3) Al   B.  X  would be the electron dot formula  1) B 2) N 3) P

  7. The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together.

  8. Electronegativity Difference • If the difference in electronegativities is between: • 1.7 to 4.0: Ionic • 0.3 to 1.7: Polar Covalent • 0.0 to 0.3: Non-Polar Covalent • Example: NaCl • Na = 0.8, Cl = 3.0 • Difference is 2.2, so • this is an ionic bond!

  9. Formation of Ions from ________ • Ionic compounds result when metals react with nonmetals • Metals loseelectrons to match the number of valence electrons of their nearest noble gas • Positive ions (cations) form when the number of electrons are less than the number of protons Group 1 metals ion 1+ Group 2 metals ion 2+ • Group 13 metals ion 3+

  10. Formation of Sodium Ion Sodium atom Sodium ion Na  – eNa + e- config: e- config: 11 p+ 11 p+ 11 e- 10 e-

  11. Formation of Magnesium Ion Magnesium atom Magnesium ion  Mg  – 2e Mg2+ e- config:e- config: 12 p+ 12 p+ 12 e- 10 e-

  12. Electron Dots For Cations • Let’s do Scandium, #21 • The electron configuration is: 1s22s22p63s23p64s23d1 • Thus, it can lose 2e- (making it 2+), or lose 3e- (making 3+) Sc = Sc2+ Sc = Sc3+ Scandium (III) ion Scandium (II) ion

  13. Electron Dots For Cations • Let’s do Silver, element #47 • Predicted configuration is: 1s22s22p63s23p64s23d104p65s24d9 • Actual configuration is: 1s22s22p63s23p64s23d104p65s14d10 Ag = Ag1+(can’t lose any more, charges of 3+ or greater are uncommon)

  14. Learning Check A. Number of valence electrons in aluminum 1) 1 e- 2) 2 e- 3) 3 e- B. Change in electrons for octet (Al) 1) lose 3e- 2) gain 3 e- 3) gain 5 e- C. Ionic charge of aluminum 1) 3- 2) 5- 3) 3+ D. 12 p+ and 10 e- 1) 0 2) 2+ 3) 2- E. 50p+ and 46 e- 1) 2+ 2) 4+ 3) 4- F. 15 p+ and 18e- 2) 3+ 2) 3- 3) 5-

  15. Ions from __________ Ions • In ionic compounds, nonmetals in 15, 16, and 17 gain electrons from metals • Nonmetal add electrons to achieve the octet arrangement (forming anions) • Nonmetal ionic charge: 3-, 2-, or 1-

  16. Chloride Ion unpaired electron octet 1 - :Cl + e  :Cl:  e- config: e- config: 17 p+ 17 p+ 17 e- 18 e- ionic charge

  17. ________ Bond • Between atoms of metals and nonmetals with very different electronegativity • Bond formed by transfer of electrons • Produce charged ions in all states. • Conductors and have high melting point. • Examples; NaCl, CaCl2, K2O

  18. Ionic bonding • Its like taking candy from a baby……. • The nonmetal (bullies) take electrons (candy) from the metals (babies)

  19. 1). Ionic bond– electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes(Na+) and the Cl becomes(Cl-), charged particles or ions.

  20. Ionic Bonding Lets do an example by combining calcium and phosphorus: • All the electrons must be accounted for, and each atom will have a noble gas configuration (which is stable). Ca P

  21. Ionic Bonding = Ca3P2 Formula Unit This is a chemical formula, which shows the kinds and numbers of atoms in the smallest representative particle of the substance. For an ionic compound, the smallest representative particle is called a: Formula Unit

  22. Properties of Ionic Compounds • ____________ solids - a regular repeating arrangement of ions in the solid: Ions are strongly bonded together. • Structure is rigid. • High melting points • Coordination number- number of ions of opposite charge surrounding it

  23. A. Energy of Bond Formation • Lattice Energy • Energy released when one mole of an ionic crystalline compound is formed from gaseous ions • Greater lattice energy = greater ionic bond C. Johannesson

  24. Do they Conduct? • Conducting electricity means allowing charges to move. • In a solid, the ions are locked in place. • Ionic solids are insulators. • When melted, the ions can move around. • Melted ionic compounds conduct. • NaCl: must get to about 800 ºC. • Dissolved in water, they also conduct (free to move in aqueous solutions)

  25. - Page 198 The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

  26. A. Oxidation Number • The charge on an ion. • Indicates the # of e- gained/lost to become stable. 1+ 0 4- 2+ 3+ 4+ 3- 2- 1- (1+ to +3)

  27. B. Ionic _________ • Write the names of both elements, cation first. • Change the anion’s ending to -ide. • Write the names of polyatomic ions. • For ions with variable oxidation #’s, write the ox. # in parentheses using Roman numerals. Overall charge = 0.

  28. Polyatomic Ions • “Many atoms” • A group of covalently-bonded atoms that have a net charge. • They act as a unit as an anion or cation. • You should be able to recognize polyatomic ions in a given chemical formula.

  29. Poly atomic ions • SymbolName • CH3COO1– acetate ion • NH41+ ammonium ion • AsO43– arsenate ion • C6H5COO1– benzoate ion • HCO31– bicarbonate ion • BrO31– bromate ion • CO32– carbonate ion • ClO31– chlorate ion • ClO21– chlorite ion • C6H5O73– citrate ion • CN1– cyanide ion • Cr2O72– dichromate ion • OH1– hydroxide ion • CrO42– chromate ion

  30. Poly atomic ions • C6H5O73– citrate ion • CN1– cyanide ion • Cr2O72– dichromate ion • OH1– hydroxide ion • ClO1– hypochlorite ion • IO31– iodate ion • PO31–phosphite ion • NO31– nitrate ion • NO21– nitrite ion • C2O42– oxalate ion • ClO41– perchlorate ion • MnO41– permanganate ion • PO43– phosphate ion • SiO32– silicate ion • SO42– sulfate ion • SO32– sulfite ion • S2O32– thiosulfate ion

  31. C. Ionic ___________ • Write each ion. Put the cation first. • Overall charge must equal zero. • If charges cancel, just write the symbols. • If not, crisscross the charges to find subscripts. • Use parentheses when more than one polyatomic ion is needed. • Roman numerals indicate the oxidation #.

  32. C. Ionic Formulas • potassium chloride • magnesium nitrate • copper(II) chloride • K+ Cl- • Mg2+ NO3-  • Cu2+ Cl-

  33. C. Ionic Formulas • calcium oxide • aluminum chlorate • iron(III) oxide • Ca2+ O2- • Al3+ ClO3-  • Fe3+ O2-

  34. B. Ionic Names • Write the names of both elements, cation first. • Change the anion’s ending to -ide. • Write the names of polyatomic ions. • For ions with variable oxidation #’s, write the ox. # in parentheses using Roman numerals. Overall charge = 0.

  35. B. Ionic Names • NaBr • Na2CO3 • FeCl3

  36. Ionic hydrates • Ionic Hydrates – ionic compounds that have loosely held water molecules • Example: CuSO4●5H20(s) Formula of raised # of water ionic compound dot molecules • By heating an ionic hydrate, the water molecules are released and the ionic compound becomes “anhydrous” Example: CuSO4 ●5H20(s) + heat CuSO4 (s) (hydrated) (anhydrous) Name: ____________________________________

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