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CHAPTER 8 AP CHEMISTRY. COVALENT BONDING  Polar bonds Uneven attraction for the shared electrons  Nonpolar bonds Bonding electrons are shared equally.

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Presentation on theme: "CHAPTER 8 AP CHEMISTRY. COVALENT BONDING  Polar bonds Uneven attraction for the shared electrons  Nonpolar bonds Bonding electrons are shared equally."— Presentation transcript:

1 CHAPTER 8 AP CHEMISTRY

2 COVALENT BONDING  Polar bonds Uneven attraction for the shared electrons  Nonpolar bonds Bonding electrons are shared equally by bonded atoms, balanced distribution of charge

3 CONTINUE MMolecule Group of two or more atoms (same or different type of elements) held together by covalent bonds (sharing electrons) and can exist in nature DDiatomic - two atoms of the same element covalently bonded CChemical formulas have symbols, numerical subscripts, and states the type of atoms and the number of each atom PPotential energy is at a minimum when attractive forces balance repulsive forces AA bond is formed if the potential energy is lower as a compound than as single atoms

4 OCTET RULE CChemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (eight) of electrons in its valence shell EExceptions to the octet rule Has less than a full octet around the central atom because the atom is too small (boron and beryllium) Has more than a full octet (10 or 12) around the central atom, central atom is large with empty d and f sublevels Central atom is the least electronegative atom Central atom is a non-metal found in periods 3, 4, and 5 and are surrounded by halogens

5 LEWIS STRUCTURE LLone pair Unshared pair of electrons, not involved in bonding and belongs to one atom Oxygen has two lone pair electrons LLewis structure FORMULAS IN WHICH ATOMIC SYMBOLS REPRESENT NUCLEI AND INNER-SHELL ELECTRONS, DOTS OR DASHES BETWEEN TWO SYMBOLS REPRESENT BONDS AND DOTS REPRESENT UNSHARED ELECTRONS

6 STRUCTURAL FORMULA AArrangement of atoms and where the bonds are formed SSingle bond: sharing two electrons DDouble bond: sharing four electrons, higher bond energy than a single bond so bond length is shorter TTriple bond: sharing six electrons, higher bond energy than a double bond, so bond length is shorter

7 BONDS  Bond energy energy required to break a bond  Ionic bonds very stable, high melting point metal and nonmetal usually form these bonds  coulomb’s law energy of interactions E = 2.31 x J. nm (Q 1 Q 2 ) r Q 1 and Q 2 numerical charge

8 CONTINUE  A bond will form when the energy in a compound is lower than as separate atoms  this will minimize the sum of positive (repulsive) energy and negative (attractive) energy terms and is called the BOND LENGTH

9 ELECTRONEGATIVITY  Ability of an atom to attract shared electrons to itself  page 334  polarity  dipole moment > б+ б- б-б- б-б- б-б- Б+Б+

10 PREDICTING IONIC COMPOUNDS  Ca + O Ca 2+ + O 2-  calcium is transferring 2 electron‘s to the oxygen  positive ions are  negative ions are  Isoelectronic ions have the same electron configuration as the noble gas it is close to.

11 LATTICE ENERGY  change in energy when gaseous ions form ionic solids  page 343  lattice energy = k(Q 1 Q 2 ), k = proportionality ( r ) constant  covalent bonds with ionic characteristics  dipole moment gives a covalent bond a percent of ionic character  (measured dipole moment X-Y) X 100 (calculated dipole moment X + Y - )  If the compound when melted can conduct an electrical current it is called a SALT

12 COVALENT BONDING  forces which hold nonmetal atoms together  because this bond consists of sharing electrons it is very stable

13 LEWIS STRUCTURES  eight electrons in the valence shell. When there are eight electrons in the outer most energy level the atom is stable  number of valence electrons = 1 to 8  In the lewis structure a covalent bond is a line.  single bond = 2 electrons shared  double bond = 4 electrons shared  triple bond = 6 electrons shared

14 WRITING LEWIS STUCTURES  draw a skeleton structure by joining the atoms by single bonds  central atom written first in the formula  the following are usually terminal atoms H, Br, F, I, Cl, O  make sure all bonds are accountable  determine the number of valence electrons that have not been accounted for

15 RESONANCE FORMS  the concept of resonance is involved whenever a single Lewis structure does not explain all bonds  Formal charge the charge on an atom in a molecule or ion that is calculated by assuming that all lone pairs of electrons shared by an atom, belong to the atom  C f = E v - (E u + 1/2E b )  C f = formal charge, E v = # of valence electrons  E u = # of unshared electrons, E b = # of bonded electrons

16 EXCEPTIONS TO THE OCTET RULE  molecules that contain an odd number of valence electrons  species that contain an odd number of electrons are paramagnetic (attracted by a charge)  less than a full octet of electrons  expanded octet surrounded by more than four pairs of valence electrons

17  enthalpy change, ΔH, to break a bond in a mole of gaseous substance  endothermic if ΔH > 0  exothermic if ΔH < 0  ΔH = (bond energies broken) -(bond energies formed) BOND ENERGY

18 CONTINUE  bond strength and length as the number of bonds increase the length decreases making then stronger  bond polarity nonpolar bond atoms share electrons equally polar bond atoms share electrons unequally ionic bond atoms do not share electrons


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