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Chemical Bonding

Chemical Bonding. Chapter 12. • The world around us is composed almost entirely of compounds. • Pure elements are rare. ✓Large numbers of the same elements bonded together: ‣ Gold nuggets ‣ Diamonds ✓Unbonded individual atoms even rarer: ‣ Argon, 1% of the atmosphere

juliette-bernard
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Chemical Bonding

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  1. Chemical Bonding Chapter 12

  2. • The world around us is composed almost entirely of compounds. • Pure elements are rare. ✓Large numbers of the same elements bonded together: ‣ Gold nuggets ‣ Diamonds ✓Unbonded individual atoms even rarer: ‣ Argon, 1% of the atmosphere ‣ He in natural gas deposits • The manner in which atoms are bound together has a profound effect on the chemical and physical properties of substances.

  3. • Diamonds and graphite are both solely composed of carbon atoms. • Diamonds are one of the hardest materials known. • Graphite is soft and slippery. • Why so different?? • The secret is in the arrangement of the atoms.

  4. • What similarities and differences do you notice between the two structures? • The thin lines between the layers are representing very loose bonds which allow the graphite layers to slip off. • Each atom in the diamond is held equally tight by four other C atoms.

  5. Chemical Bonding • What is a chemical bond? ✓Remember: protons are + and electrons are - ✓Two or more atoms coming together and maximizing their +/- attractions while minimizing their -/- and +/+ repulsions. ✓This balancing of interactions occurs when electrons are concentrated between two nuclei. ✓Voila! This is a bond • How do chemists learn about bonds? ✓By measuring the bond energy - the energy needed to break a bond. ✓By observing which atoms, and in what ratios bond with each other

  6. Types of Bonding • Metallic Bonding - swirling valence e- ✓The “sea of valence electrons” surrounding the array of positive ions (nucleus + inner core of e-). • Ionic bonds - transfer of e- ✓Metals bonding with nonmetals resulting in +/- ions that attract each other. • Covalent (or Molecular) bonds - sharing e- ✓Nonpolar Covalent ‣ When the same (or similar) nonmetal atoms bond together ‣ Equal sharing of electrons ✓Polar covalent bonds ‣ When different nonmetal atoms bond ‣ Unequal sharing of electrons

  7. How to Decide if its Equal or Not?Electronegativity • A number that tells us the ability of an atom (engaged in a bond) to attract electrons to itself. • Polarity of a bond depends on the difference between the electronegativities of two atoms. ✓(nonpolar) ≤ 0.4 < polar < 2.0 ≤ (ionic) • Any molecule that has a positive end and a negative end is said to be ✓Polar ✓aka: a dipole ✓aka: has a dipole moment.

  8. Lewis Structures • A schematic drawing to show the arrangement of the valence electrons in a molecule, used to predict valid atom arrangements • Each atom is trying to “satisfy” their electronic “desire” for an octet. (Hydrogen only “needs” a duet.) • All valence electrons must be accounted for in the Lewis structure. • Each electron is represented by a dot . ✓Each pair represented by two dots .. ✓Or each pair represented by a line —

  9. Electronic Nirvana • Each element is “searching” for an octet ✓8 valence electrons ✓Except for hydrogen which only needs 2 • We will consider ✓Bonding pairs ‣ Electrons shared between two nuclei ✓Nonbonding, lone pairs, or unshared pairs ‣ Electrons not involved in bonding

  10. Drawing Lewis Structures • Add up the valence electrons in all the atoms in the molecule. • Assemble the framework by attaching each atom with a single bond. • Complete the octets of all the other atoms with nonbonding pairs. • If there are not enough electrons to go around ✓“stretch” the electrons by getting “double duty” out of them by making one or more (if necessary) multiple bonds • If writing the structure of a polyatomic ion, the negative charge indicates the number of electrons that must be added to the total valance electrons.

  11. Rules that will help predict an accurate framework. • Oxygen atoms do not like to stick to each other except in the molecule O2 , O3 , or in peroxide molecules. • H’s and F’s are always terminal atoms, on the outside – attached to only one other atom. • H’s and F’s only form single bonds. (They do not form double or triple bonds.) • Carbon atoms love to hook together with each other and do not like to be terminal atoms (on the outside), because they do not like to have unshared pairs of electrons. If at all possible, avoid putting a carbon in a position where it needs unshared pairs. • Sometimes chemists write the chemical formula in the order that the atoms attach to each other. ✓ Use the order of atoms as a guide, not as a hard and fast rule. ✓ Often the first atom listed is the central atom. ✓ More electronegative

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