Chapters 7 and 8

1. Chapters 7 and 8 Chemical bonding

2. Valence Electrons • Elements with similar chemical behavior have the same number of valence electrons. • For the representative elements (1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A) the group number corresponds to the number of valence electron in each group (with the exception of He) • When examining electron configurations, the electrons that are present in the highest principle energy level represent the valence electrons of those atoms. Br: [Ar]4s23d10 4p5 Bromine has 7 valence electrons

3. Valence Electrons and Electron Dot Structures • Valence electrons are the electrons that participate in chemical bonds • Electron dot structures consist of the atom symbol and its valence electrons represented as dots. Br: [Ar]4s2 3d10 4p5

4. The Octet Rule - Metals • Metals tend to lose electrons to gain stability. When they lose electrons they attain a stable electron configuration where the next lowest-energy level has a full octet.

5. The Octet Rule – Non-metals • Non-metals tend to gain electrons to obtain a full octet in their highest occupied p sublevel

6. Formation of Ionic Compounds • Cations and anions combine to form ionic compounds NaCl + +

7. The Formation of Calcium Chloride CaCl2

8. Properties of Ionic Compounds • The chemical formula for an ionic compound is called a formula unit which represents the lowest who number ration of atoms in that compound.( ex. NaCl) • Most ionic compounds are crystalline solids at room temperature • Ionic compounds generally have very high melting points • Ionic compounds can conduct electricity when molten or dissolved in water Coordination Number = The number of ions of opposite charge surrounding the ion in a crystal.

9. Metallic Bonding • Positive metal ions floating in a sea of electrons. • Metals are excellent conductors of electricity • Metals can be easily deformed • (They are malleable and ductile) Know that metals are arranged in crystalline structures, but do not worry about the arrangement of these structures: face-centered, body-centered, hexagonal.

10. Alloys • Alloys are mixtures of metals • Properties of alloys are usually superior to those of their component elements. Interstitial Alloy Substitutional Alloy

11. Covalent Bonding • Covalent bonds occur between two or more non-metals • Unlike ionic bonds where electrons are transferred from one atom to another, electrons are shared between atoms in a covalent bond. • Atoms joined together by covalent bonds are called molecules • A compound composed of molecules is called a molecular compound

12. Molecular and Structural Formulas • A molecular formula indicates the types and numbers of each atom in a molecule • The structural formula indicate the arrangement of the atoms in the molecule H2O

13. Covalent Bonds and the Octet Rule • Atoms share electrons in a covalent bond so that each atom has enough electrons to satisfy the octet rule

14. Varieties of Covalent Bonds • Single bonds (sigma bonds) en • One pair of electrons is shared between two atoms Lone pair

15. Varieties of Covalent Bonds • Double Bonds (1sigma bond, 1 pi bond) • Atoms share two pairs of electrons • Triple Bonds (1sigma bond, 2 pi bonds) • Atoms share three pairs of electrons

16. Coordinate Covalent Bonds • A covalent bond in which one atom contributes both bonding electrons.

17. Resonance Structures Def: A condition when more than one valid Lewis structure can be written for a molecule or ion.

18. Exceptions to the Octet Rule Too few electrons surrounding the central atom (ex: BH3) Boron will not have a full octet, only 6 electrons. It can only achieve a full octet when another atom shares an entire pair of electrons with it (Coordinate covalent bonding) Too many electrons surrounding the central atom (ex: PCl5) An odd number of electrons

19. How to Draw a Lewis Structure for Molecules Predict the location of atoms If there are more than two atoms, place the least electronegative atom in the center and surround it by the remaining atoms. Hydrogen is always terminal (outside) because it can only make one bond Determine the total number of electrons if each atom had a full set of valence electrons (2 for H, 8 for all others) Add up the number of valence electron that you have to work with Subtract total valence electrons from total electrons and divide by two. This is the number of bonding pairs that are needed to put together the molecule. Connect the atoms with the number of bonds that you calculated above Add lone pairs where needed so that each atom has a full octet (except for hydrogen which can only have two electrons)

20. Polyatomic Ions Polyatomic ions are a cluster of non-metals that carry a charge. To draw the structure of a polyatomic ion, follow the procedure for drawing ordinary molecules but add or subtract the number of electrons gained or lost to the total number of valence electrons in your structure as indicated by the charge on the ion.

21. Molecular Shape (VSEPR) Valence Shell Electron Pair Repulsion – minimizes the repulsion of shared and unshared pairs of electrons around the central atom. The shape of a molecule determines many of its physical and chemical properties. The VSEPR is based on the arrangement of bonding and lone electrons around a central atom to minimize repulsion. The repulsion of electrons creates a specific bond angle between a central atom and two terminal atoms. Lone pairs of electrons occupy more space than bonding pairs of electrons

22. Molecular Geometry or 2

23. Electronegativity and Polarity Recall: Electronegativity is the ability of an atom to attract an electron.

24. Chemical bonding is like “Tug-o-War” Electronegativity Bond Type Difference Non-polar Covalent 0-0.4 Polar Covalent 0.5-2.0 Ionic >2.0

25. Molecular Polarity Molecules are either polar or non-polar Both polar and non-polar molecules may contain polar bonds. What determines whether a molecule is polar or non-polar is the symmetry of the molecule VSEPR shapes that can demonstrate symmetry are: Linear Trigonal Planar Tetrahedral