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. When atoms release much energy as a result of bond formation, the bond is considered to be strong and stable. Thus in order to break the bond, the substance must absorb much energy. When atoms release much energy as a result of bond formation, the bond is considered to be strong and stable. Thus in order to break the bond, the substance must absorb much energy. When atoms release little energy as a result of bond formation, the bond is considered to be weak and unstable. Thus in order to break the bond, the substance must absorb little energy. When atoms release little energy as a result of bond formation, the bond is considered to be weak and unstable. Thus in order to break the bond, the substance must absorb little energy.

6. 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.

7. 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

8. Representation of an Ionic Bonding [Na] + [Cl] - [Cl] - [Mg] 2+ [Cl] -

10. c) Covalent Bonds: Characteristics 1. Are bonds between nonmetals. 2. Involve the equal and unequal sharing of electrons. 3. Are considered to be relatively weak bonds. * * Exceptions are network solids (diamonds, graphite, silicon dioxide, asbestos, silicon carbide)

12. 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

13. 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).

14. Note: Most atoms will share a single pair of electrons that is represented by a single bond (a dash). Most atoms will share a single pair of electrons that is represented by a single bond (a dash).

15. Other atoms have the ability to share two or three pairs of electrons representing a double bond (two dashes) or a triple bond (three dashes). Other atoms have the ability to share two or three pairs of electrons representing a double bond (two dashes) or a triple bond (three dashes). O OO OO OO O ::.. Double bond (4 electrons shared) N N :: Triple bond (6 electrons shared)

16. 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.

17. 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.

19. 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).

22. 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.

23. 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.

24. 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).

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

26. Hydrogen bonds between water molecules

27. 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.

28. 1. Molecular Polarity and Symmetry If a compound has non-polar bonds (as with the diatomic molecules), the molecule is considered to be non-polar. If a compound has non-polar bonds (as with the diatomic molecules), the molecule is considered to be non-polar. Polar and Nonpolar Molecules: A polar molecule is one that has an uneven electron distribution resulting in a net positive and negative charge in different parts of the compound. A polar molecule is one that has an uneven electron distribution resulting in a net positive and negative charge in different parts of the compound. A non-polar molecule is one that has an even distribution of electrons resulting in no net charge. A non-polar molecule is one that has an even distribution of electrons resulting in no net charge.

29. In most cases, if a molecule contains a polar covalent bond, it is considered to be a polar molecule. In most cases, if a molecule contains a polar covalent bond, it is considered to be a polar molecule. The exception exists when the molecule is completely symmetrical. The exception exists when the molecule is completely symmetrical. A symmetrical molecule will have an even distribution electrons resulting in no net charge. A symmetrical molecule will have an even distribution electrons resulting in no net charge. Methane (CH 4 ): Tetrahedral Structure Carbon Dioxide (CO 2 ): Linear Structure

30. The four molecules that have polar covalent bonds but are non-polar molecules due to their symmetry are: CH 4 - Methane CCl 4 - Carbon Tetrachloride CF 4 - Carbon Tetrafluoride CO 2 - Carbon Dioxide

31. b) Van der 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.

32. 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.

33. Types of Bonds in Compounds with Polyatomic Ions Compounds with two different element are called binary compounds. The types of bonds found between atoms in a binary compound are dependent on the types of elements. Compounds with two different element are called binary compounds. The types of bonds found between atoms in a binary compound are dependent on the types of elements. Examples of such compounds include: CO 2 NaClCaF 2 Li 2 O Carbon Dioxide Sodium Chloride Calcium Fluoride Lithium Oxide Compounds with three or more different elements are called ternary compounds. Compounds with three or more different elements are called ternary compounds. If a ternary compound has a metal and a polyatomic ion (see reference table E), it will have both ionic and covalent bonds. If a ternary compound has a metal and a polyatomic ion (see reference table E), it will have both ionic and covalent bonds.

34. Na 2 SO 4 KNO 3 Sodium Sulfate Potassium Nitrate Li 3 PO 4 Lithium Phosphate Mg(OH) 2 Magnesium Hydroxide

35. Covalent Bond (bond between two non-metals) Ionic Bond (bond between a metal and a non-metal)

36. VII. Formulas a)Chemical Formula: represents the type and number of element(s) in a chemical compound. The type of element is represented by its chemical symbol. The type of element is represented by its chemical symbol. The number of atoms of the element in the compound is represented by its subscript.The number of atoms of the element in the compound is represented by its subscript. Li 3 PO 4 Lithium = Li Phosphorus = P Oxygen = O 3 atoms 1 atoms 4 atoms NOTE: Subscripts are the numbers found to the lower right of a given chemical symbol. If a subscript is not present assume that it is a number one.

37. Types of Formulas 1.Molecular Formulas Represents the type and number of elements in a covalent compound.Represents the type and number of elements in a covalent compound. The subscripts do not need to be reduced in a molecular formula.The subscripts do not need to be reduced in a molecular formula. Ex: C 2 H 4 Ethene C 6 H 12 O 6 Glucose C 6 H 12 O 6 Glucose

38. 2. Empirical Formulas Represents the lowest whole number ratio of atoms in a compound. Represents the lowest whole number ratio of atoms in a compound. For ionic compounds, the formula is always empirical. This is due to the fact that one must reduce the subscripts in an ionic compound to the lowest whole number ratio. For ionic compounds, the formula is always empirical. This is due to the fact that one must reduce the subscripts in an ionic compound to the lowest whole number ratio. Ex: Calcium Sulfide Ca S Ca +2 S -2 Ca 2 S 2 Crisscross Method Disregard the signs. Ca S Reduce

39. For covalent compounds (molecules), reducing its molecular formula to an empirical formula only shows the ratio of elements within the compound. For covalent compounds (molecules), reducing its molecular formula to an empirical formula only shows the ratio of elements within the compound. Ex: C 6 H 12 O 6 Molecular Formula C 1 H 2 O 1 Empirical Formula Ratio: C:H:O = 1:2:1

40. H2OH2OH2OH2O Molecular and Empirical Formula

41. 3. Structural Formulas Structural formulas indicate the type, number, and arrangement of atoms within a compound. Structural formulas indicate the type, number, and arrangement of atoms within a compound. O H H O = C = O [Na] + [Cl] -

42. Glucose

43. VIII. Formula Writing a) Ionic compounds: Given the name of the compound: 1. Determine if the compound is binary or ternary. 2. If the compound is binary, write out the symbols for the two elements present: metal then non-metal. * Elements with lower electronegativity values are usually listed first in a chemical formula. 3. Find and write out the oxidation number for the metal and non-metal. 4. Perform the crisscross method to determine the subscripts. * Eliminate the signs from the subscripts and if necessary reduce the subscripts to the lower whole number.

45. 5.If the compound is ternary, refers to Reference Table E. 6.Write out the symbol for the metal and the polyatomic ion. 7.Place parenthesis around the polyatomic ion. 8.Write in the oxidation number for the metal and the ionic charge of the polyatomic ion. 9. Perform the crisscross method to determine the subscripts. * Eliminate the signs from the subscripts and if necessary reduce the subscripts to the lower whole number.

46. Naming Ionic Compounds with a Transition Metals Elements that can have more than one possible charge MUST have a Roman Numeral to indicate the charge on the individual ion. Example #1 : Copper (II) SulfateCuSO 4 Copper’s oxidation number is +2. Sulfate’s charge is -2. Sulfate’s charge is -2. The Roman Numeral “ II ” indicates the oxidation number of the transition metal. Example #2: Iron (III) OxideFe 2 O 3 Iron’s oxidation number is +3. Oxygen’s oxidation number is -2. Oxygen’s oxidation number is -2. The Roman Numeral “III” indicates the oxidation number of the transition metal.

47. Examples of Older Names of Cations formed from Transition Metals

48. Complete the names of the following binary compounds with variable metal ions: FeBr 2 iron (_____) bromide CuClcopper (_____) chloride SnO 2 ___(_____ ) ______________ Fe 2 O 3 ________________________ Hg 2 S________________________

49. Complete the formulas of the following ternary compounds with variable metal ions: 1. copper(II) nitrate 2.Iron (III) hydroxide 3. Tin(IV) hydroxide

50. b) Covalent Compounds (molecular substances) Prefix System (binary compounds) 1. 1.Less electronegative non-metal comes first. 2. Add prefixes to indicate # of atoms. Omit mono- prefix on the FIRST element. Mono- is OPTIONAL on the SECOND element. 3. Change the ending of the second element to -ide.

51. Greek Numerical Prefixes 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca

52. Given the names of the following molecules, determine their formulas. Dinitrogen monoxide Potassium sulfide Dichlorine heptoxide Iodine monochloride

53. Given the formulas of the following molecules, determine their names. BaI 2 P4S3P4S3P4S3P4S3 I2O5I2O5I2O5I2O5 CS 2 B 2 Cl 4