Chemical Bonding Reference Table - Ionic and Covalent Bonds

1. PACKET #7: Chemical Bonding Reference Table: PT & Table S www.regentsprep.org

2. Chemical Bonding • Chemical Bond: an attraction between the protons of one atom and the electrons of the next atom that attaches the atom together. • Formed by transferring or sharing of electrons. • A chemical bond has stored or potential energy. • After a chemical bond is formed, the atoms have a complete outer shell they are stable.

3. Bonds can be classified as being either polaror non-polar. • Polarity: tendency of a molecule, or compound, to be attracted or repelled by electrical charges because of an asymmetrical arrangement of atoms around the nucleus. • Think of it like a game of tug of war, if one end of the compound is pulling on the electrons more than the other, there is an unequal pull, and therefore, the substance is polar. If there is an equal pull, then the substance is non-polar. • This concept of polarity is determined by electronegativity.

4. Electronegativity Difference Remember . . . • Electronegativity: an atom’s attraction for electrons in a bond. • The higher the EN, the more the atom attracts electrons. • The lower the EN, the less the weaker the attraction for electrons. • REFER TO TABLE S: Note which elements have higher EN (TRENDS).

5. Ionic Bond • Attraction between oppositely charged ions • Occurs when electrons are transferred from one ion (charged particle) to another • Electronegativity difference 1.7+ • Metals react with Nonmetals to form ionic compounds • Always Polar !!!

6. Forming Ionic Bonds • Ionic bonds are formed when valence electrons are transferredfrom metals to non-metals forming ionic compounds.  • Metals lose electrons and become cations (+). • Nonmetals gain electrons and become anions (-) • Example: KCl:

7. Example: CaBr2 • An ionic bond was formed by the TRANSFER OF ELECTRONS!! (NO SHARING)!!! • Draw the following ionic compounds: SrF2, LiI, BaCl2, Na2O, AlCl3, and Al2O3

8. Polyatomic Ions • Ionic compounds containing polyatomic ions can have both ionic and covalent bonding. • Example: KNO3 • Notice that NO3- is composed of 2 nonmetals therefore the bonding is covalent between N and O but the bonding between K+ and NO3-  is ionic.

9. Example: (NH4)3PO4 • The bonding is covalent between N & H, as well as between P & O, but ionic between the two individual polyatomic ions. • Draw the Lewis structure of the following ionic compounds NaOH, Mg(NO3)2 and (NH4)2CO3

10. Possible Combinations for Ionic Compounds

11. Properties of Ionic Compounds • Hard • Good conductors of electricity in liquid or aqueous form only, because ions can move in solution and in liquid form, but not in solid form. • High melting and boiling points • Solid at room temperature • Dissolve in polar substances: like water. (Polar – opposite charges).

12. Covalent Bonds • Formed when 2 atoms (both nonmetals) shareelectrons. [Example Cl2 or H2O] • Neither atom pulls strongly enough to remove an electron from the other • The EN difference is < 1.7 • Unpaired electrons pair up in such a way that the atoms complete their outer shells • Covalent compounds also referred to as molecular compounds

13. Properties of Covalent Bonds • Gases, liquids or solids • Soft • Nonmetals • Poor conductors of heat and electricity because they are not charged particles. (No ions or mobile electrons) • Low melting and boiling points because of weak attraction between molecules.

14. Polar vs. Non-Polar Covalent Bonds • Unlike an ionic compound, a covalent compound can be classified as either a polar covalent bond, or a non-polar covalent bond. • If the EN of the atoms are different then it is a polar covalent bond. • If the EN of the atoms are the same or very similar then it is a non-polar covalent bond. 0.0 - 0.4 = non-polar covalent 0.5 - 1.6 = polar covalent

15. Polar Covalent Bonds • There will always be an unequal sharing of electrons due to the EN difference. • Example: HCl • EN of H = 2.1 EN of Cl = 3.2 • Difference is 1.1, which is less than 1.7, but greater than 0.4

16. Non-Polar Covalent Bonds • EN difference 0.0 – 0.4 • The non-polar covalent bonds you must commit to memory are the diatomic molecules: • H2, N2, O2, F2, Cl2, Br2, I2 • These are considered covalent bonds because they are two non-metals sharing electron, and are considered non-polar because since they are the same element, they have the same EN, and therefore the difference is 0. • Since “likes dissolve in likes” non-polar covalent compounds will only dissolve in non-polar solvents.

17. Number of Covalent Bonds • Single covalent bond: one pair of shared electrons; 2 electrons total • -Double covalent bond: two pairs of shared electrons; 4 electrons total • -Triple covalent bonds: three pairs of shared electrons; 6 electrons total

18. A Little Review . . . The Octet Rule: • Atoms seek to have eight valence electrons when in a bond. • All Noble Gases have a full octet. • Exceptions to the octet rule:H and He will only hold 2 electrons each. He already has 2 electrons and will not bond with other atoms.

19. Rules for Drawing Lewis Diagrams of Covalent Compounds • Calculate the total number of valence electrons available to the molecule or ion. 2. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by outer atoms. Hydrogen isNEVERthe central atom. 3. Form bonds between the central atom and outer atoms with a pair of electrons. All remaining electrons should be distributed so that each atom has 8 electrons.

20. Example #1 - Methane CH4 1. Determine the total number of valence electrons available: • One carbon has 4 valence electrons.Four hydrogen, each with one valence electron, totals 4. • This means there are 8 valence electrons, making 4 pairs, available.

21. Example #2 - Ammonia NH3 1. Determine the total number of valence electrons available: • One nitrogen has 5 valence electrons.Three hydrogen, each with one valence electron, totals 3. • This means there are 8 valence electrons, making 4 pairs, available.

22. Example #3 - Carbon tetrachloride CCl4 1. Determine the total number of valence electrons available: • One carbon has 4 valence electrons.Four chlorine, each with 7 valence electrons, totals 28. • This means there are 32 valence electrons, making 16 pairs, available.

23. Example #4 – O2 & N2 DOUBLE BONDTRIPLE BOND

24. MORE  . . . • Some other covalent compounds that are important to know how to draw are: H2, F2, Cl2, Br2, I2, O2, N2, HCl, HF, HBr, HI, CO2, CF4, CBr4, H2O, H2S, CCI4. SiH4

25. Partially Positive & Negative • In a polar covalent bond, both of the elements are non-metals, and therefore there is no “true” + or – charges; instead there are partially (+) and partially (-) charges. • The element with the higher EN is partially (-) and the one with the lower EN is partially (+)

26. Polar & Non-PolarBondsvs. Polar & Non-PolarMolecules • Just to add a little more confusion  . . . • RECALL: Polar Bond = different EN Non-Polar Bond = same (similar) EN BUT Polar Molecule = Asymmetrical Non-Polar Molecule = Symmetrical

27. Polar Molecule • Polar molecules are asymmetrical (the compound is not a mirror image if you folded it over itself) • Polar molecules result from an unequal distribution of electrons. • These molecules are also called dipoles.

28. Polar Molecules

29. Non-Polar Molecules • Non-polar molecules have an equal distribution of electrons throughout the entire compound. • The electrons are being pulled in all directions evenly. • All diatomic molecules are symmetrical

30. This is a SNAP! • Symmetric are • Nonpolar • Asymmetric are • Polar

31. Let’s Practice!! • For the following covalent bonds, determine there bond type and molecule type (EXPLAIN WHY!!!) • H2O • N2 • NH3 • CO2 • CH4

32. HF H2O CO2 CH4 NH3 HCl H2S CS2 CF4 PH3 HBr SiO2 CCl4 HI SiS2 CBr4 CI4

33. Exceptions to the Octet Rule Less than an Octet: • When there are fewer than eight electrons around an atom in a molecule or ion. • This is a rare situation and is most often encountered in compounds of boron and beryllium • Example: BF3

34. Exceptions to the Octet Rule More than an Octet: • When there are more than eight electrons in the valence shell of an atom. Much more common than having less than eight. • Examples: PCl5 (10 valence electrons) ICl4- (12 valence electrons)

35. Other Types of Covalent Bonds Coordinate Covalent Bond: • When one atom donates both of the electrons that are shared • Example: NH4+ and H3O+ • Nitrogen donates a pair of electrons to share with H+ forming a coordinate covalent bond between nitrogen and hydrogen

36. Other Types of Covalent Bonds Network Solids: • Solids that have covalent bonds between atoms linked in one big network or one big macromolecule with no discrete particles. This gives them some different properties from most covalent compounds. • They are hard, poor conductors of heat and electricity, and have high melting points • Examples include: Diamond (C), silicon carbide (SiC), and silicon dioxide (SiO2)

37. Metallic Bond • Occurs only in metals (Example Copper) • Metals have low ionization energy meaning they hold onto their valence electrons very loosely • As a result the electrons in metallic substances move about very easily and are not associated with any particular atom • Therefore, the particles of a metal are usually positive ions surrounded by a mobile sea of electrons • The attraction between the positive cations and the moving electrons is what holds the metal together • Properties of Metallic Bonds are that of metals: hard, good conductors of heat & electricity, malleable, ductile, etc . . .

38. Intermolecular Forces • Forces of attraction between molecules. Include: dipoles, hydrogen bonds, dispersion forces, and molecule-ion attraction. • The difference between intra- (within) and inter- (between). A covalent bond would be an intra-molecular force (a bond within a molecule). Dipole-Dipole attractions would be a inter-molecular force (bonds found between molecules) • The higher the degree of polarity in the bonds the stronger the intermolecular forces

39. Dipole-Dipole Attractions • Positive end of a polar molecule is attracted to the negative end of an adjacent polar molecule.

40. Hydrogen Bonding • An intermolecular attraction between a hydrogen atom in one molecule to a nitrogen, oxygen, or fluorine atom in another molecule • The strongest intermolecular force • Substances with hydrogen bonds tend to have much higher melting and boiling points than those without hydrogen bonds • Example: The boiling point of H2O is much higher than H2S

41. Hydrogen Bond

42. London Dispersion ForcesAKA: van der Waals Forces • Weak intermolecular forces between non-polar molecules (like diatomic molecules) • Dispersion forces make it possible for small, non-polar molecules to exist in both liquid or solid phases under conditions of high or low temperatures. • Increases with molecular size, Ex. As you go down group 17, dispersion forces increase and boiling point increases.

43. Molecule-Ion Attraction • Attraction between the ions of an ionic compound such as NaCl, and a molecule such as water (or any other polar covalent compound). • When you put NaCl into water, the Na+ from the salt is attracted to the O from the water which is partially (-), and the Cl- from the salt is attracted to the H+ of the water.

44. Molecular Geometry • Lewis dot structures do notshow us what the shape of a molecule is. • Molecular geometry allows us to learn the relationship between the two-dimensional Lewis dot structures that we learned to draw, and the three-dimensional molecular shapes that we are about to learn about. CCl4

45. Molecular Shapes • There are five fundamental shapes.

46. Review Questions 1) Which type of bond results when one or more valence electrons are transferred from one atom to another? 1) a nonpolar covalent bond 2) a polar covalent bond 3) a hydrogen bond 4) an ionic bond 2) Which compound contains ionic bonds? 1) CO2 2) CaO 3) NO2 4) NO

47. 3) Which molecule contains a triple covalent bond? 1) N2 2) H2 3) Cl2 4) O2 4) What is the total number of electrons shared in the bonds between the two carbon atoms in a molecule of 1) 6 2) 2 3) 8 4) 3 5) Covalent bonds are formed when electrons are 1) mobile within a metal 2) transferred from one atom to another 3) captured by the nucleus 4) shared between two atoms

48. 6) Which molecule contains a nonpolar covalent bond? 7) Which formula represents a nonpolar molecule? 1) CH4 2) H2S 3) HCl 4) NH3

49. 8) Which pair of characteristics describes the molecule illustrated below? 1) symmetrical and polar 2) asymmetrical and nonpolar 3) symmetrical and nonpolar 4) asymmetrical and polar 9) Which intermolecular force of attraction accounts for the relatively high boiling point of water? 1) ionic bonding 2) covalent bonding 3) metallic bonding 4) hydrogen bonding

50. 10) In the diagram of an ammonium ion below, why is bond A considered to be coordinate covalent? 1) Nitrogen provides a pair of electrons to be shared with hydrogen. 2) Hydrogen provides a pair of electrons to be shared with nitrogen. 3) Nitrogen transfers a pair of electrons to hydrogen. 4) Hydrogen transfers a pair of electrons to nitrogen. 11) Which structural formula represents a polar molecule?