1. Chemical Bonding Unit IV

2. I. Chemical Bonds: are attractive forces that hold atoms and/or compounds together. result from the simultaneous attraction of an atom’s positively charged nucleus for other atoms negatively charged electrons. result when electrons are transferred (ionic) and shared (covalent) between atoms. results in the increase in the chemical stability of atoms when energy is released (exothermic).

3. II. Chemical Bonds and Energy a) To break a chemical bond, the given chemical must absorb energy from the environment. BOND BREAKING IS AN ENDOTHERMIC PROCESS. b) When chemical bonds are formed, the stability of the reactants is generally increased. The reactants stability is increased as they release energy. BOND MAKING IS AN EXOTHERMIC PROCESS. NOTE: WHEN CHEMICALS CONTAIN LARGE AMOUNTS OF ENERGY, THEY ARE CONSIDERED TO BE UNSTABLE and are VERY REACTIVE. WHEN CHEMICALS CONTAIN SMALL AMOUNTS OF ENERGY, THEY ARE CONSIDERED TO BE STABLE and are UNREACTIVE.

4. III. Electronegativity: a) Definition: Refers to the force of attraction that a given atom has for another atom’s valence electron(s). The higher the electronegativity, the greater the atoms force of attraction on electrons. b) Characteristics : 2. Nonmetals have greater electronegativity values than metals. (Nonmetals have a greater tendency to attract electrons as compared to metals.) 1. Down a group (top to bottom) – electronegativity decreases. The lower the electronegativity, the lesser the atoms force of attraction on electrons. 1. Based upon the atom of Fluorine, which has the highest electronegativity value of 4. c) Periodic Trends 2. Across a period (left to right) – electronegativity increases with increasing nuclear charge (# of protons).

5. IV. Types of Bonds Intramolecular Forces – bonds between atoms a) Ionic Bonds: Characteristics 1. Ionic bonds exist between ions (charged particles). 2. Ionic bonds involve the transfer of electrons (one atom gains and the other loses). 3. Ionic bonds are generally formed between a metal (electron donor) and a nonmetal (electron acceptor). 4. Ionic bonds are considered to be very strong bonds. 5. Ionic bonds are predicted by a difference in electronegativity greater than 1.7.

6. b) Ionic Compounds: Characteristics 1. High melting and boiling points. 2. Solids at STP (standard temperature and pressure). 3. Form crystal lattice structures. 4. Conduct electricity in the liquid and aqueous phases only. 5. Have regular geometric arrangements. 6. 6.Examples include: NaCl (Sodium Chloride) LiCl (Lithium Chloride) KF (Potassium Fluoride) MgBr 2 (Magnesium Bromide) MetalNonmetal Free ions present

8. c) Covalent Bonds: Characteristics 1. Are bonds between nonmetals. 2. Involve the equal and unequal sharing of electrons. 3. 3.Are considered to be relatively weak bonds 4.Low melting point and low boiling point 5.Mainly found in liquid and gas state, some solids 6.Poor conductors of heat and electricity no free ions (exception – Acids are conductors in aqueous solution)

9. 4. TYPES: polar covalent bonds – unequal sharing of electrons; occurs between two different nonmetals; electronegativity difference is between 0 and 1.7. ex: H 2 O, SO 2, CH 4 Sulfur dioxide molecule

10. nonpolar covalent bonds - equal sharing of electrons; occurs between two of the same nonmetals; electronegativity difference is equal to 0. ex: H 2, O 2, N 2, Cl 2, Br 2, I 2, F 2 - DIATOMS coordinate covalent bonds – sharing of electrons whereas both of the electrons of the shared pair are donated by the same atom (not required for the regents exam).

11. 5. Characteristics of Covalent Compounds Are also known as molecular substances. Exist in all three phases at STP. Poor conductors of heat and electricity (good insulators). Have low melting and boiling points.

12. Diamond Graphite Network Solids (covalent compounds) NOTE: Network Solids are covalent compounds that are extremely hard and have very high melting and boiling points. These represent exceptions to the general rules of covalent compounds. Examples include: diamonds, graphite, SiO2, and SiC.

14. d) Metallic Bonds - very strong bonds between metal ions. “positive ions immersed in a sea of mobile electrons”. Bonds between metals. ex: Ag(s), Mg(s), Ca(s) Metallic Compounds are lustrous, ductile, malleable, and excellent conductors of heat and electricity. Solids at STP (except Mercury (Hg) which is a liquid).

17. V. The Octet Rule Electrons are found outside of the nucleus of atoms. Electrons are arranged around the atom’s nucleus according to the amount of energy they possess. Electrons are found in energy levels. The outermost energy level (valence shell) for a given atom contains valence electrons. In the case of most atoms of the periodic table, for the atom to be stable, it must contain a total of 8 electrons in its valence level. To become stable atoms will gain, lose, or share electrons to obtain this valence shell configuration.

18. Exceptions to the Octet Rule include: hydrogen, helium, lithium, beryllium, sulfur, nitrogen. More precisely, for an atom to be stable, it must have an electron configuration similar to that of a noble gas.

19. VI. Bonds Between Molecules a) Dipole-Dipole Interactions Covalent compounds that have unequal sharing of electrons (polar covalent bonds) are considered to be polar molecules. The unequal sharing of electrons results in molecules that have positively charged and negatively charged regions. The positive region of one polar molecule will be attracted to the negative region of another polar molecule (vice versa). Intermolecular Forces that are weaker in comparison to intramolecular forces (ionic, covalent, and metallic bonds).

20. From: http://www.geo.arizona.edu/xtal/geos306/9_7.jpg

21. Hydrogen bonds between water molecules

22. Hydrogen bonds. An example of dipole-dipole interactions. Oxygen has a greater electronegativity than hydrogen. Thus, the electrons are shared unequally. The shared electrons are drawn closer to the oxygen atoms giving the them a slight negative charge. The hydrogen atoms have a slight positive charge. The negatively charged oxygen of one water molecule is attracted to the positively charged hydrogen of the other water molecule.

23. b) Van deer Waals Forces Weak forces of attraction between nonpolar compounds. Forces of attraction between: Monoatomic Molecules (Noble Gases) Diatomic Molecules Other select nonpolar substances. Remember, most nonpolar compounds contain nonpolar bonds which equally distribute electrons between the atoms.

24. c) Molecule-Ion Attractions Forces of attraction between ions and polar covalent compounds. Example: Which of the following contain molecule-ion attractions? 1. CaCl 2 (s) 2. CO 2 (g) 3. NaCl (aq) 4. Ag (s) Answer: Choice 3. Sodium chloride (being a salt, ionic compound, and an electrolyte) will dissolve, dissociate, and form ions in water. Water is a polar covalent molecule. Thus…. Molecule- Ion Attraction. Hint: Always look for the choice that has an ionic compound in the aqueous state.