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Covalent Bonding

Covalent Bonding. (Continued). Electronegativity. Crack out table S, pp. 10-11 in the reference tables The modern definition of electronegativity is due to Linus Pauling. It is the power of an atom in a molecule to attract electrons to itself .

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Covalent Bonding

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  1. Covalent Bonding (Continued)

  2. Electronegativity • Crack out table S, pp. 10-11 in the reference tables • The modern definition of electronegativity is due to Linus Pauling. It is the power of an atom in a molecule to attract electrons to itself. • SOMEWHERE in the your tables, jot this down: • Electronegativitydifference determines bond type. • 0.0 <__> .9 – non-polar covalent • 1.0 <__> 1.6 – polar covalent • 1.7 or more – ionic

  3. Polarity in Bonding(Pt 1) • The more electronegative element holds on to the electrons more tightly and this creates a slightly negative charge on that end of the molecule. • The other end of the molecule becomes slightly positively charged. • This creates a Dipole and the whole molecule becomes like a magnet with a positive end and a negative end.

  4. Single Covalent Bonds • Atoms held together by sharingone pair of electrons are joined by a covalent bond • Hydrogen gas consists of diatomic molecules(H2) whose atoms share only one pair of electrons, forming a single covalent bond. • diatomic molecule- a molecule consisting of two atoms: 7 main diatomic molecules: • N2, O2, H2 and halogens, group 17: • F2, Cl2, Br2I2

  5. Hydrogen • An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots. • The pair of shared electrons forming the covalent bond is also often represented as a dash, as in H—H for hydrogen. • A structural formula represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms.

  6. Fluorine • Because a fluorine atom has seven valence electrons, it needs one more. • By sharing electrons and forming a single covalent bond, two fluorine atoms each achieve the electron configuration of neon. • The two fluorine atoms share one pair of valence electrons. A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair or a nonbonding pair

  7. Water • Water (H2O) is a molecule containing three atoms with two single covalent bonds. • Two hydrogen atoms share electrons with one oxygen atom. • In the electron dot structures, the oxygen atom in water has two unshared pairs of valence electrons. • Draw:

  8. Methane • Methane (CH4) contains four single covalent bonds. • The carbon atom has four valence electrons and needs four more. • Each of the four hydrogen atoms contributes one electron to share with the carbon atom, forming four identical carbon-hydrogen bonds. • In the electron dot structure, methane has no unshared pairs of electrons. • Draw:

  9. Double and Triple Covalent Bonds • Sometimes atoms bond by sharing more than one pair of electrons. • Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons. • A bond that involves two shared pairs of electrons is a double covalent bondaka, a double bond • A bond formed by sharing three pairs of electrons is a triple covalent bond aka, a triple bond

  10. The Carbon Problem • Carbon can have FOUR bonding sites and can share more than one pair of electrons. • Draw:

  11. Polarity in Bonding (pt 2) Hydrogen Bonding • This attraction between the hydrogen of one water molecule and the oxygen of another water molecule is strong compared to other dipole interactions. • This relatively strong attraction, which is also found in hydrogen-containing molecules other than water, (“NOF” about hydrogen bonds) is given the special name hydrogen bond.

  12. H-Bonds Continued • The H atom is a proton and an electron. • When the electron is in a bond, the proton is left exposed. • This exposed proton will attract itself to any electron in the area. • If the electron it finds is tightly held (atom with high electronegativity) a “tug of war” results that hold the molecules together tighter.

  13. Water’s H-Bonding • Hydrogen bonds are the strongest of the intermolecular forces. • They are extremely important in determining the properties of water and biological molecules such as proteins. • The relatively strong attractive forces between water molecules causes the water to form small drops on a waxy surface.

  14. Draw

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