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Chemical Shapes & Polarity

Chemical Shapes & Polarity. Here are some basic compounds that you see everyday. Do you know what they are? As you can see, even though they are made of the same elements (carbon and hydrogen), they have very different shapes. This is due to the way the atoms bond to each other.

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Chemical Shapes & Polarity

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  1. Chemical Shapes & Polarity

  2. Here are some basic compounds that you see everyday. Do you know what they are? As you can see, even though they are made of the same elements (carbon and hydrogen), they have very different shapes. This is due to the way the atoms bond to each other.

  3. In this chart you can see the 6 basic shapes of molecules. Most molecules in the world take one or more of these shapes. This chart is nice because it shows the number of valence electrons involved in bonding.

  4. Here are a few of the odd molecule shapes you may encounter. The figure shows the non-bonding pairs of electrons. Remember on a Lewis Dot Structure, you may have only a single dot or a double dot. The double dots do not bind to anything. Those double dots are non-bonding pairs. Non-bonding electrons Na O Bonding electrons

  5. As seen in the figure to the left, each shape has a particular bond angle. This angle is determined by how much stress in placed on the bonds. Electrons do not like to be near each other and will push apart as much as possible. Bond angles do not change. Look at the trigonal planar molecule. The electrons can only push so far before the electrons on the other side begin to push back. These forces keep the bonds in the same place.

  6. Practice building the following molecules using the chemical models and draw stick models: NaOH P2O5 CaCO3 (NH4) 3PO4 CH3COOH NH4OH H2SO4 MgNH4PO4 AlCl3 Ba(OH) 2

  7. When two identical atoms form a covalent bond, as in H2 or Cl2, each have equal share of the electron pair. the electron density is the same on both ends of the bond, because the electrons are equally attracted to both nuclei.

  8. But when different atoms bond, as in HCl, one nucleus attracts the electrons from a bond more strongly than the other. The electrons are shared, but unequally. There are partial charges on both sides of the bond, indicated by a lowercase Greek letter delta, . In HCl, the chlorine carries a partially negative charge (-) while the hydrogen carries a partially positive charge (+). This is because the chlorine atom attracts the electrons stronger than the hydrogen.

  9. Polarity A bond that carries a partial charge is called a polar covalent bond. Because there are two poles of charge involve, the bond is a dipole. If the entire molecule has a partial charge, it is a polar molecule. This relative attraction of electrons in a bond is called electronegativity. In HCl, the Cl is more electronegative than the H and attracts the electrons. The electrons spend more time around the Cl than the H. Looking it up on a periodic table or a similar table, one can see the H has an electronegativity of 2.1 while Cl has one of 2.9. Therefore, the electronegativity difference is 0.8. Electronegativity decreases as you go down a group and increases as you go left to right in a period.

  10. Because the bonding pair in the carbo-fluorine bond is pulled towards the fluorine end of the bond, that end is left rather more negative than it would otherwise be. The carbon end is left rather short of electrons and so becomes slightly positive. The symbols + and  - mean "slightly positive" and "slightly negative". You read  + as "delta plus" or "delta positive". We describe a bond having one end slightly positive and the other end slightly negative as being polar.

  11. The polar nature of the elements sets up an overall charge on the molecule. This overall charge is represented by the symbol and called a dipole moment. Each dipole moment for each bond faces the - molecule. Then all the dipole moments are “added” and face the general direction of all the other dipole moments.

  12. These 2 molecules have individual dipole moments for each bond, but they face opposite directions, so they cancel each other out for a dipole moment of 0.

  13. Polarity & Shape

  14. Intermolecular Forces • Secondary bonds, called intermolecular forces, are much weaker than primary bonds (covalent, ionic, or metallic). They often provide a "weak link" for deformation or fracture. Example for secondary bonds are: • Hydrogen Bonds • Van der Waals Bonds

  15. A molecule’s polarity will cause it to be attracted to + and - ions or other polar molecules. When a nonpolar molecule is exposed to moving water, it won’t move with the water. So, when a polar molecule exposed to moving water, the polar molecule will move too.

  16. Hydrogen bonds are common in covalently bonded molecules which contain hydrogen, such as water (H2O). Since the bonds are primarily covalent, the electrons are shared between the hydrogen and oxygen atoms. However, the electrons tend to spend more time around the oxygen atom. This leads to a small positive charge around the hydrogen atoms, and a negative charge around the oxygen atom. When other molecules with this type of charge transfer are nearby, the negatively charged end of one molecule will be weakly attracted to the positively charged end of the other molecule. The attraction is weak because the charge transfer is small.

  17. Van der Waals bonds are very weak compared to other types of bonds. The electrons surrounding an atom are always moving. At any given point in time, the electrons may be slightly shifted to one side of an atom, giving that side a very small negative charge. This may cause an attraction to a slightly positively charged atom nearby, creating a very weak bond. Van Der Waals bonding is a that exists between virtually all atoms or molecules, but its presence may be obscured if any of the three primary bonding types is present. Secondary bonding forces arise from atomic or molecular dipoles. In essence, an electron dipole exists whenever there is some separation of positive and negative portions of an atom or molecule. When an electron cloud density occurs at one side of an atom or molecule during the electron flight about the nucleus, Van Der Waals forces are generated. This creates a dipole wherein one side of the atom becomes electrically charged and the other side has deficiency of electrons and is considerably charged positive. The atom is distorted as shown in the Figure.

  18. Which molecules are polar? Nonpolar? NaOH P2O5 CaCO3 (NH4) 3PO4 CH3COOH NH4OH H2SO4 MgNH4PO4 AlCl3 Ba(OH) 2

  19. Let’s do some practice! Find the molecular weight of: Find the weight of: NaOH 1. 5 moles of NaOH HCOOH 2. 1.2 moles of HCOOH Al2(SO4)3 3. 0.005 moles of Al2(SO4)3 H3PO4- 4. 10 moles of H3PO4-

  20. More practice with moles: Suppose you have 25 g of NaCl. How many moles do you have? Suppose you have 0.5 g of P2O5. How many moles do you have? Suppose you have 9.41 grams of Al2(SO4)3. How many moles do you have?

  21. Draw the Lewis Dot Structures for: NaOH P2O5 CaCO3 (NH4) 3PO4 CH3COOH NH4OH H2SO4 MgNH4PO4 AlCl3 Ba(OH) 2

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