1. Chem I Chapter 6 Chemical Bonding Notes

2. Chemical Bond – a mutual attraction between the nuclei and valence electrons of different atoms that binds the atoms together. Chemical Bond – a mutual attraction between the nuclei and valence electrons of different atoms that binds the atoms together.

3. Types of Chemical Bonds Ionic bond – force of attraction between oppositely charged ions, which are formed by the TRANSFER of electron(s) from metal atoms to nonmetal atoms. Ionic bond – force of attraction between oppositely charged ions, which are formed by the TRANSFER of electron(s) from metal atoms to nonmetal atoms. –Metals LOSE electrons easily! –Nonmetals GAIN electrons easily!

4. Covalent Bond – Bond resulting from SHARING electron pairs between two atoms. (Nonmetal to nonmetal) Covalent Bond – Bond resulting from SHARING electron pairs between two atoms. (Nonmetal to nonmetal) Metallic Bond – Bond between metal atoms involving a cluster of cations (Positive charged ions) surrounded by a “sea of electrons” Metallic Bond – Bond between metal atoms involving a cluster of cations (Positive charged ions) surrounded by a “sea of electrons”

5. Covalent bonds Nonpolar covalent bond – bond electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. (Hydrogen – Hydrogen bond) Nonpolar covalent bond – bond electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. (Hydrogen – Hydrogen bond) Polar covalent bond – bond in which the bonded atoms have an unequal attraction for the shared electrons. (water is an example) Polar covalent bond – bond in which the bonded atoms have an unequal attraction for the shared electrons. (water is an example)

6. Use electronegativity differences to predict whether a bond is ionic, polar-covalent, or nonpolar covalent. Use electronegativity differences to predict whether a bond is ionic, polar-covalent, or nonpolar covalent. Use pg. 151 Table to find the electronegativities for each element and Pg162 to determine bonding type. Use pg. 151 Table to find the electronegativities for each element and Pg162 to determine bonding type. Example – Cl bonded to Ca Example – Cl bonded to Ca Cl= 3, Ca=1 3-1 = 2 which makes it ionic Cl= 3, Ca=1 3-1 = 2 which makes it ionic

7. 6-2 Covalent Bonding and Molecular compounds Molecule- a neutral group of atoms that are held together by covalent bonds. Molecule- a neutral group of atoms that are held together by covalent bonds. Molecular compound- a chemical compound whose simplest units are molecules. Molecular compound- a chemical compound whose simplest units are molecules. Chemical formula – indicates the relative #’s of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. Chemical formula – indicates the relative #’s of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. Molecular formula – shows the types and #’s of atoms combined in a single molecule of a molecular compound Molecular formula – shows the types and #’s of atoms combined in a single molecule of a molecular compound

8. Atoms in molecules are joined by covalent bonds. Atoms in molecules are joined by covalent bonds. Bond length – The distance between two bonded atoms at their minimum potential energy. Bond length – The distance between two bonded atoms at their minimum potential energy. Bond energy – the energy required to break a chemical bond and form neutral isolated atoms. Bond energy – the energy required to break a chemical bond and form neutral isolated atoms.

9. The shorter the bond length the stronger the bond. The shorter the bond length the stronger the bond. The longer the bond length the weaker the bond. The longer the bond length the weaker the bond. See bond lengths pg. 168 Table 6-1 See bond lengths pg. 168 Table 6-1

10. Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (8) electrons in its highest occupied energy level (valence shell) Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (8) electrons in its highest occupied energy level (valence shell) Exceptions to the octet rule: H – needs only 2 electrons, Boron (B) – needs only 6 electrons. Exceptions to the octet rule: H – needs only 2 electrons, Boron (B) – needs only 6 electrons. Expanded Valences – when combining with Fluorine, oxygen, and chlorine - other elements may have more than 8 electrons because it involves electrons in d orbitals as well as in s and p. Expanded Valences – when combining with Fluorine, oxygen, and chlorine - other elements may have more than 8 electrons because it involves electrons in d orbitals as well as in s and p.

11. Electron –dot notation: an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. Electron –dot notation: an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. See pg. 170 Figure 6-10 for examples. See pg. 170 Figure 6-10 for examples. Unshared pair – is a pair of electrons that is not involved in bonding and belongs exclusively to one atom. Unshared pair – is a pair of electrons that is not involved in bonding and belongs exclusively to one atom.

12. Lewis Structures Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons. Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons.

13. Structural formula – indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule. Ex. F-F, H-Cl Structural formula – indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule. Ex. F-F, H-Cl

14. Single and multiple bonds Single bond – is a covalent bond produced by the sharing of one pair of electrons between two atoms Single bond – is a covalent bond produced by the sharing of one pair of electrons between two atoms Double bond – sharing of two pairs of electrons Double bond – sharing of two pairs of electrons Triple bond – sharing of three pairs of electrons. Triple bond – sharing of three pairs of electrons. Double and Triple bonds are referred to as multiple bonds. Double and Triple bonds are referred to as multiple bonds. Triple bonds are the strongest – single bonds are the weakest. Triple bonds have the shortest bond length. Triple bonds are the strongest – single bonds are the weakest. Triple bonds have the shortest bond length.

15. Resonance Structure Some molecules and ions cannot be represented adequately by a single Lewis structure. These molecules or ions have an average of two or three structures. Together these structures are refrered to as RESONANCE structures or hybrids. To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures. Some molecules and ions cannot be represented adequately by a single Lewis structure. These molecules or ions have an average of two or three structures. Together these structures are refrered to as RESONANCE structures or hybrids. To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures.

16. 6.3 Ionic Bonding & Ionic compounds Ionic compound – composed of positive and negative ions that are combined so that the total (or overall) charge is zero. Ionic compound – composed of positive and negative ions that are combined so that the total (or overall) charge is zero. An ionic compound is usually a crystalline solid made up of a three-dimensional arrangement of ions called a CRYSTAL LATTICE. An ionic compound is usually a crystalline solid made up of a three-dimensional arrangement of ions called a CRYSTAL LATTICE. Ionic compounds are not, therefore, composed of independent units (molecules) Ionic compounds are not, therefore, composed of independent units (molecules)

17. The formula for an ionic compound is not a molecular formula. Instead it shows only the simplest ratio of element in the compound. The formula for an ionic compound is not a molecular formula. Instead it shows only the simplest ratio of element in the compound. Empirical Formula-Simplest ratio of elements in a compound Empirical Formula-Simplest ratio of elements in a compound Formula Unit – Amount of an ionic compound represented by its formula. Formula Unit – Amount of an ionic compound represented by its formula.

18. Lattice Energy – Amount of energy released when one mole of an ionic crystal is formed from gaseous ions. Lattice Energy – Amount of energy released when one mole of an ionic crystal is formed from gaseous ions.

19. Comparison of Ionic & Molecular compounds Ionic Ionic –High melting points and boiling points –Hard, brittle crystalline solids –Most are soluble in water –Conduct electricity in molten state or in solution

20. Molecular Molecular –Lower melting and boiling points –Most are liquids or gases –Polar molecules soluble in water – nonpolar are not –Generally non conductors of electricity

21. Polyatomic Ions – groups of atoms which are covalently bonded together and which carry an overall charge Polyatomic Ions – groups of atoms which are covalently bonded together and which carry an overall charge Resonance structures – refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure Resonance structures – refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure Structural formulas - Chemical formulas which show all information contained in lewis structure EXCEPT unshared electrons. Structural formulas - Chemical formulas which show all information contained in lewis structure EXCEPT unshared electrons.

22. 6.4 Metallic bonding Metals lose electrons easily Metals lose electrons easily Metals contain few valence electrons and, therefore, have any empty orbitals. Metals contain few valence electrons and, therefore, have any empty orbitals. These empty orbitals overlap in a metallic crystal, allowing valence electrons to pass freely from atom to atom. These are referred to as DELOCALIZED electrons. These empty orbitals overlap in a metallic crystal, allowing valence electrons to pass freely from atom to atom. These are referred to as DELOCALIZED electrons. Metal atoms form a lattice of cations surrounded by a “sea of electrons” Metal atoms form a lattice of cations surrounded by a “sea of electrons”

23. Metallic Properties Good conductors of heat & electricity Good conductors of heat & electricity Shiny Shiny Malleability – ability to be hammered into sheets Malleability – ability to be hammered into sheets Ductility – ability to be stretched into a wire Ductility – ability to be stretched into a wire

24. 6.5 Molecular Geometry VSEPR Theory – Valence Shell Electron Pair Repulsion VSEPR Theory – Valence Shell Electron Pair Repulsion States that electron pairs in the valence shell of an atom repel each other and, therefore, arrange themselves around the atom as far apart as possible. States that electron pairs in the valence shell of an atom repel each other and, therefore, arrange themselves around the atom as far apart as possible.

25. Electron Pair Geometry Number of electron pairs on the central atom determines electron pair geometry Number of electron pairs on the central atom determines electron pair geometry Arrangement of atoms determines MOLECULAR GEOMETRY Arrangement of atoms determines MOLECULAR GEOMETRY Double and Triple bonds should be considered 1 pair of electrons when determining molecular geometry. Double and Triple bonds should be considered 1 pair of electrons when determining molecular geometry.

26. Bond Angles A Bond angle is the angle between two bonds involving a single atom. A Bond angle is the angle between two bonds involving a single atom. See Chapter 6 Section 5 for specific angles See Chapter 6 Section 5 for specific angles

27. Intermolecular forces Forces of attraction between molecules are intermolecular forces Forces of attraction between molecules are intermolecular forces Boiling point is a good measure of intermolecular forces, since these forces must be overcome during boiling – the higher the boiling point the stronger the intermolecular force. Boiling point is a good measure of intermolecular forces, since these forces must be overcome during boiling – the higher the boiling point the stronger the intermolecular force. Intermolecular forces are typically weaker than bonds in metals, Ionic compounds and molecules. (see chart, pg. 190) Intermolecular forces are typically weaker than bonds in metals, Ionic compounds and molecules. (see chart, pg. 190)

28. Kinds of I-M attractions Dipole – equal but opposite charges separated by a short distance. Dipole – equal but opposite charges separated by a short distance. An arrow with a + sign at the positive end and the arrow head pointing to the negative end shows the charge direction. An arrow with a + sign at the positive end and the arrow head pointing to the negative end shows the charge direction. Dipole-Dipole Forces – force of attraction between two polar molecules. Dipole-Dipole Forces – force of attraction between two polar molecules.

29. Dipole – Induced Dipole Force is an attraction between a polar molecule and a nonpolar molecule. Dipole – Induced Dipole Force is an attraction between a polar molecule and a nonpolar molecule. Hydrogen bonding – Strong I-M force of attraction between a hydrogen atom bonded to a VERY electronegative atom (N, O, F) in a nearby molecule – strongest of I-M forces. Hydrogen bonding – Strong I-M force of attraction between a hydrogen atom bonded to a VERY electronegative atom (N, O, F) in a nearby molecule – strongest of I-M forces.

30. London Dispersion Forces – I-M force of attraction between nonpolar molecules which results from the constantly shifting positions of electrons. The weakest of the I-M forces. London Dispersion Forces – I-M force of attraction between nonpolar molecules which results from the constantly shifting positions of electrons. The weakest of the I-M forces.

31. Strength of I-M forces Hydrogen BondingStrongest Hydrogen BondingStrongest Dipole-Dipole Dipole-Dipole Dipole- Induced Dipole Dipole- Induced Dipole London Dispersion Weakest London Dispersion Weakest

32. Polarity of Molecules Polaritiy of a molecule depends upon: Polaritiy of a molecule depends upon: –Polarity of bonds, –Orientation of bonds

33. Hybridization The “mixing” of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies. The “mixing” of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies. The type of hybridization which occurs is indicated by arrangement of electron pairs in the molecule. The type of hybridization is designated by the number and kind of orbitals taking part in the process The type of hybridization which occurs is indicated by arrangement of electron pairs in the molecule. The type of hybridization is designated by the number and kind of orbitals taking part in the process