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Covalent Bonds and Compounds

Covalent Bonds and Compounds. Molecules. Three Kinds of Bonds. Non-metal to non-metal metal to non-metal metal to metal. Covalent Ionic Metallic. Bonds and Electronegativity. Electrons are transferred between atoms when the difference in electronegativity between the atoms is quite high.

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Covalent Bonds and Compounds

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  1. Covalent Bonds and Compounds Molecules

  2. Three Kinds of Bonds • Non-metal to non-metal • metal to non-metal • metal to metal • Covalent • Ionic • Metallic

  3. Bonds and Electronegativity • Electrons are transferred between atoms when the difference in electronegativity between the atoms is quite high. • The amount of transfer depends on the electronegativity difference.

  4. Bonds and Electronegativity • The number 1.67 seems to be the magic number.(Note:1.67 rounds to 1.7) • If the electronegativity difference is less than 1.67, the bond is more covalent than ionic. • If the electronegativity difference is greater than 1.67, the bond is more ionic than covalent.

  5. Electronegativity Difference • Covalent < 1.67 However – • 0 -.6 is non-polar covalent • . 6 – 1.67 polar covalent • There are 7 instances of perfectly covalent bonds (electronegativity difference = 0) • H2, N2, O2, F2, Cl2, Br2, I2

  6. Covalent bonds are generally between two non-metals. CO CO2 Carbon monoxide Carbon dioxide Nomenclature (naming)

  7. 1 2 3 4 5 6 7 8 9 Mon(o) Di Tri Tetr(a) Pent(a) Hex(a) Hept(a) Oct Non Prefixes

  8. N2O NO N2O3 NO2 N2O5 Dinitrogen monoxide Nitrogen monoxide Dinitrogen trioxide Nitrogen dioxide Dinitrogen pentoxide Nitrogen and oxygen(five molecules)

  9. Organic Molecules(hydrocarbons) • Composed primarily of carbons (always) and hydrogens (usually). • Three primary types • Alkanes – only single bonds • Alkenes – at least one double bond • Alkynes – at least one triple bond

  10. Root name from # of Carbons • 1 – meth • 2 – eth • 3 – prop • 4 – but • 5 – pent • 6 – hex • 7 – hept • 8 – oct • 9 – non • 10 - dec

  11. H C- H-C- C- C- C- C- C- C- H-C-H C- H H H H H H - - Some examples methane CH4 # of carbons ethane C2H6 Notice: # of H is 2n+2 C3H8 Propane butane C4H10

  12. ethane H-C- C- H H H H-C=C-H H H H H Alkenes ethene # of H’s is just 2n

  13. H H H H H H H H H ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ H-C-C=C-C-C-C-C-C-C-H H-C-C-C=C-C-C-C=C-C-H H-C-C-C-C-C-C-C-C-C-H ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ ɩ H H H H H H H H H H H H H H H H ɩ ɩ ɩ H H H H H H H H H H H H H H H H H H H H H H H Saturated vs. Unsaturated 2-nonene nonane 2,6 nonadiene

  14. Formula Mass • Mass of one formula unit. • Add the mass of all of the elements times their subscripts. • Sodium phosphate: • Na3PO4 • Na = 22.99 x 3 = 68.97 amu • P = 30.97 x1 = 30.97 amu • O = 16.00 x 4 = 64.00 amu • 68.97 + 30.97 + 64 = 163.94 amu

  15. Percent Composition • What is the percent of sodium in sodium phosphate? • What is the mass of sodium in Na3PO4? • Na = 22.99 x 3 = 68.97 amu • What is the mass of Na3PO4? • 163.94 amu • % of Na is 68.97/163.94 x 100 or • 40.07%

  16. Empirical Formulas • A formula in lowest terms. (Use the GCF) • Greatest Common Factor • The empirical formula for C2H6 is • CH3 • What is the formula mass for C2H6? • What is the formula mass for CH3? • What is 30 ÷ 15?

  17. Determining Empirical Formula from Percent composition. • A certain compound contains 32.38% sodium, 22.65% sulfur, and 44.99% oxygen. Find the empirical formula. • We’re going to do percent composition in reverse… • Assume a 100g sample – therefore 32.38g of sodium etc.

  18. 0.7063 0.7063 0.7063 32.28g 1 mole 22.99g = 1.408 moles = 2 22.65g 1 mole 32.07g = 0.7063 moles Smallest value = 1 44.99g 1 mole 16.0g These are the subscripts = 2.812moles = 4 Divide all values by the smallest of them(this will give whole numbers) Determine the molar ratios • So, the empirical formula is Na2SO4 Na S O

  19. Actual Formula from Empirical Formula • If you know the empirical formula and the formula mass of the actual formula, you can determine actual formula by finding the formula mass of the empirical formula and dividing it into the formula mass of the actual formula. This will give you the GCF. • Multiply the subscripts of the empirical formula by the GCF.

  20. Example • What is the formula of a compound whose molecular mass is 150.1amu and its empirical formula is CH2O? • Formula mass of CH2O is 30.02amu. • 150.1 ÷ 30.02 = 5 (That’s the GCF) • Multiply the subscripts 1,2,1 by the GCF (5) which gives the new subscripts of 5,10,5 or • C5H10O5

  21. Sea of “delocalized” electrons Metal atoms Metallic Bonds • Bonding in metals is due to delocalized electrons. • These often exist in what is called a sea of electrons.

  22. Metallic Bonds • This explains many of the properties of metals: • Malleable • Ductile • Conducts electricity well

  23. Alloys • Two Metals (and sometimes other substances) bonded (mixed) together.

  24. Alloys • Two Metals (and sometimes other substances) bonded (mixed) together.

  25. Alloys • Two Metals (and sometimes other substances) bonded (mixed) together.

  26. Alloys • Two Metals (and sometimes other substances) bonded (mixed) together.

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