Chapter 4 – Compounds and Their Bonds

1. Chapter 4 – Compounds and Their Bonds • Study Goals • Octet Rule and Ions •  Ionic Compounds •  Naming and Writing Ionic Formulas •  Polyatomic Ions •  Covalent Compounds •  Electronegativity and Bond Polarity • Shapes and Polarity of Molecules

2. The Bondings in Compounds In nature, almost all elements in the Periodic Table form compounds having 2 or more elements combined together. A compound is a pure substance consisting of 2 or more different types of atoms combined together (e.g. CO2, NaCl). In this chapter, we are going to learn about the types of interaction or bonding that holds atoms together in a compound. Several types of molecular bondings are known that hold atoms together in a compound. The 2 most fundamental types of bonding we are going to discuss is the ionic bonding and the covalent bonding. COVALENT BONDING Bonding occurs as a result of electrostatic attraction between oppositely charged particles. Bonding occurs as a result of sharing of a pair of electrons between 2 atoms.

3. Ionic Compounds Ionic compounds include many everyday's substances that we are familiar with – table salt (NaCl), baking soda (NaHCO3), milk of magnesia (Mg(OH)2) and calcium carbonate (CaCO3) used as antacids. Ionic compounds also include ionic crystals such as those found in naturally occurring gems, including ruby (Al2O3) and sapphires. - ionic compounds are generically called “salts”. NaCl Ruby: Al2O3 (Impurity: Chromiun) NaHCO3 CaCO3 Sapphire: Al2O3 (Impurities: Iron, Titanium)

4. Covalent compounds are more predominant than ionic compounds and are made of elements that are non-metals. They include substances that we already know such as water (H2O), table sugar (sucrose = C12H22O11), propane gas (C3H8), and drugs such as the antibiotic amoxicillin (C16H19N3O5S) and the antidepressant Prozac (C17H18F3NO). - covalent compounds are referred to as “molecules”. Covalent Compounds Water (H2O) Table sugar (sucrose) Propane gas (C3H8) Prozac (C17H18F3NO )

5. Recall that the valence shell e− are electrons of the s- and p-orbitals – these are the electrons of the outermost shell and they are the ones that usually participate in chemical reactions. • the outermost shell, when filled with a complete octet of valence electrons, provide a particularly stable electronic configuration to the atom. Natural elements that have a complete octet of electrons are the noble gases. • - when elements combine to form compounds – whether ionic or covalent compounds – they usually do so to acquire this octet configuration of electrons around them so that they achieve the most stable chemical state─ this is the key to understanding chemical reactions during which atoms combine to form compounds. Formation of Positively Charged Ions – Cations ● Many atoms can ionize (i.e. lose or gain extra e–) and acquire a net charge in the process. The tendency to lose or gain e– is a characteristic of a particular group of elements.

6. Group 1A elements (alkali metals): • have 1 valence e– in the outermost shell. • tend to lose this e– to achieve the noble gas configuration • provided by the filled lower shell; as it loses one e–, a • group 1A element acquires a +1 charge. • hence, all group 1A elements lose 1 e– to form ions • with a +1 charge. • Ex: Na  Na+ + e– • Li  Li+ + e–

7. Group 2A elements (alkaline earth metals): • have 2 valence e– in the outermost shell. • tend to lose these 2 e– to achieve the noble gas • configuration provided by the filled lower shell. As it • loses 2 e–, a group 2A element acquires a +2 charge. • hence, all group 2A elements lose 2 e– to form ions • with a +2 charge. • Ex: Mg  Mg2+ + 2 e– • Ca  Ca2+ + 2 e– In general: A metal tends to lose its outermost shell electron(s) to become a positively charged ion (cation) when it ionizes. In doing so, it achieves the stable electronic configuration of a noble gas.

8. Group 3A elements: • have 3 valence e– in the outermost shell. • tend to lose these 3 e– to achieve the noble gas • configuration. • as it loses 3 e–, a group 3A element acquires a +3 charge. • Ex: Al  Al3+ + 3 e– • - similarly pattern of ionization is seen for the other group 3A elements: Gallium (Ga), Indium (In), and Thallium (Tl). Al  Al3+ + 3 e– Ga  Ga3+ + 3 e– In  In3+ + 3 e– Tl  Tl3+ + 3 e– An exception: Although Boronis a group 3A element, it does NOT form a 3+ ion like the elements below it (e.g. Al, Ga, In, and Tl) because its outer shell electrons are much closer to the (+)charged nucleus and therefore much more difficult to be removed. As one goes down the column, however, the increasingly larger outer shells put the outermost electrons further away from the nucleus, thus allowing some of these electrons to be removed more easily. In addition, the e–’s in the outermost shells are also partially shielded from the nucleus with layers of electrons from the inner shells. Hence, while Boron does not form B3+ ion, all the other elements within group 3A do.

9. PERIODIC TABLE

10.  There is a tendency for some elements of groups 3A  5A to lose only the p-electrons when they ionize. For examples: TlTl+ (loss of 1 p-electron); Tl3+ (loss of 1 p- and 2 s-electrons) SnSn2+ (loss of 2 p-electrons); Sn4+ (loss of 2 s- and 2 p-electrons) PbPb2+ (loss of 2 p-electrons); Pb4+ (loss of 2 s- and 2 p-electrons) Bi Bi3+ (loss of 3 p-electrons) - note that these cations are formed only from elements whose valence e− are at higher energy levels and therefore are much further removed from the nucleus. They are therefore the exception and generally not the rule, for the reason that cations with increasingly larger (+) charges are harder to form due to increasingly larger ionization energies that are required to remove the electrons. Ti3+ Ti4+ Se2– Te2– Tl+ Tl3+ Bi3+

11. Formation of Negatively Charged Ions – Anions • Group 7A elements (the Halogens): • - have 7 valence e– in the outermost shell. • tend to gain 1 e– to achieve the noble gas • configuration. • as it gains 1 e–, a group 7A element acquires a • –1 charge. Hence, all group 7A elements gain 1 e– • to form ions with a –1 charge. Negatively charged ions are called anions. • Ex: F + 1 e– F– (fluoride ion) • Cl + 1 e– Cl– (chloride ion)

12. Group 6A elements: • have 6 valence e– in the outermost shell. • tend to gain 2 e– to achieve the noble gas • configuration, thus acquiring a −2 charge. • hence, all group 6A elements gain 2 e– to form • ions with a –2 charge. • Ex: O + 2 e– O–2 (oxide ion) • S + 2 e– S–2 (sulfide ion)

13. Group 5A elements: • have 5 valence e– in the outermost shell • tend to gain 3 e– to achieve the noble gas • configuration, thus acquiring a −3 charge. • Ex: N + 3 e– N–3 (nitride ion) • P + 3 e– P–3 (phosphide ion) • - interestingly, whereas nitrogen (N) and phosphorus (P) gain 3 e– to achieve the octet configuration as they ionize, antimony (Sb), and bismuth (Bi) lose e–upon ionization. This is most likely due to the fact that antimony and bismuth are further removed from the nucleus, thereby their outer e– are more easily removed upon ionization. By contrast, the outer shell e– of nitrogen and phosphorus are closer to the nucleus and therefore not so easily removed due to stronger nuclear attraction. Instead, these atoms gain e– to achieve the octet configuration as they ionize. In general: A nonmetal tends to gain electrons in its outer shell to become a negatively charged ion (anion) when it ionizes. In doing so, it achieves the stable electronic configuration (octet configuration) of a noble gas.

14. The transition metals also form positive ions, and many have variable valence (i.e. having more than one oxidation state). Ex: Cr2+ Cr3+ Fe2+ Fe3+ Cu+ Cu2+ Au+ Au3+ Pb2+ Pb4+ Sn2+ Sn4+ - the (2+) ions form as a result of loss of the 2 s-orbital electrons. Higher oxidation state results from additional loss of d-orbital electrons (e.g. Fe3+ = loss of two s-orbital e– and one d-orbital e–). - some transition metal elements form only one type of cation: Ag+, Cd2+, Zn2+ Transition Metals

15. Names of Ions • Metal ionsform onlypositive ions and are named the same way as their • elemental names: • Ex: Na+ sodium ion • Li+ lithium ion • Ca2+ calcium ion • Al3+ aluminum ion • For metal ions that have variable charge, a Roman numeral that indicates the charge of the ion is placed in parentheses following the name of the metal: Metal ion IUPAC name Common name Cr2+ chromium(II) chromous ion Cr3+ chromium(III) chromic ion Cu+ copper(I) cuprous Cu2+ copper(II) cupric Fe2+ iron(II) ferrous Fe3+ iron(III) ferric

16. Metal ion IUPAC name Common name Pb2+ lead(II) plombous Pb4+ lead(IV) plombic Sn2+ tin(II) stannous Sn4+ tin(IV) stannic Co2+ cobalt(II) cobaltous Co3+ cobalt(III) cobaltic Au+ gold(I) Au3+ gold(III) • Nonmetal ions are usually negatively charged and are named by replacing the ending of their elemental name with -ide: Nonmetal ion Name of ion N3– nitride P3– phosphide O2– oxide S2– sulfide F– fluoride

17. When metals react with nonmetals, e– are transferred from the metal atoms to the nonmetal atoms, forming ions. The resulting compound is called an ionic compound. Examples: Sodium (Na) + Chlorine (Cl)  Sodium Chloride (NaCl)Magnesium (Mg) + 2 Chlorine (Cl)  Magnesium Chloride (MgCl2) - in each of these reactions, the metal atom tranfers its outer shell e– to the nonmetal atom, so that the metal atom becomes positive ion and the nonmetal atom becomes negative ion. There is a strong electrostatic force of attraction between these oppositely charged ions, called an ionic bond. As a result of e– transfer, ions are formed and each has a noble gas configuration. Formation of Ionic Compounds Ex: the Na atom has lost its single 3s e– and taken on a [Ne] electronic configuration, whereas the Chlorine has accepted one e– into its 3p subshell and taken on the [Ar] configuration. - ions with such noble gas configurations are particularly stable, accounting in part for the stability of the formation of the ionic crystal NaCl.

18. Structure of the Sodium Chloride Crystal Within the NaCl crystal, each Na+ ion is surrounded by six Cl– ions, and likewise each Cl– ion is surrounded by six Na+ ions. - in an ionic crystal made of Na+ and Cl– ions, the notion of a well-defined NaCl molecule does not exist (contrast this to a water (H2O) molecule, a covalent compound). At the molecular level, the crystal is a regular alternating array of Na+ ions and Cl– ions. However, because the ratio of Na+ ions to Cl– ions is 1:1, we say that the formula of sodium chloride consists of one sodium atom and one chlorine atom, NaCl. NaCl crystal

19. 2 Na + S  Na2S charge balance: 2 Na+ : 1 S–2 • The subscript 2 indicates 2 Na+ ions are needed to balance the charge of 1 S–2 ion. • - because a metal tends to lose e– to acquire (+)charge when it forms ion and a non-metal gains e– to acquire (–)charge when it forms ion, ionic compounds are made from the combining of a metal element with a nonmetal element. • the attractive force in an ionic compound is purely ionic– that is, 2 different elements form a compound as a result of their opposite charges attracting one another. An ionic compound is also referred to generically as a salt. • - Other examples of Ionic Compounds: BaF2 MgCl2 LiBr Ionic Compound: Sodium Sulfide

20. Ionic Compound: Magnesium Chloride Mg + 2 Cl  MgCl2charge balance: 1 Mg2+ : 2 Cl– The subscript 2 indicates 2 Cl– ions are needed to balance the charge of 1 Mg+2 ion.

21. Naming Ionic Compounds An ionic compound is named first with the metal ion which is followed by the name of the nonmetal ionending in -ide: Ex: NaF  sodium fluoride MgBr2 magnesium bromide Al2O3 aluminum oxide - note that the subscripts are NOT mentioned in naming ionic compounds. The subscripts are understood to be placed in the formula correctly in order to balance the charge of the ions: NaF  Na+ + F– MgBr2 Mg2+ + 2 Br– Al2O3 2 Al3+ + 3 O2– Natural red sapphire made of Al2O3. Color is due to presence of transition metal impurities. Sodium fluoride (NaF) is an active ingredient in toothpaste, preventing the formation of cavities.

22. ● When an ionic compound consists of a metal ion with variable charge, the charge of the metal cation can be determined from the nonmetal anion. The metal ion is then named using a Roman numeral placed in parentheses to indicate its charge: Ex: Name the compounds Cu2S and SnO2. Cu2S: Since the nonmetal sulfide ion (S–2) carries a –2 charge, the metal copper ion must be Cu+1 because there are 2 copper atoms in the formula. The name of the compound is then copper(I) sulfide (formal IUPAC name). This compound is also known by the common name cuprous sulfide. Chalcocite - a copper sulfide ore SnO2: The nonmetal oxide ion (O–2) has a –2 charge. Therefore the charge on the tin metal ion is +4. The compound is named tin(IV) oxide. Cassiterite is a mineral that contains SnO2 and is an important source of tin.

23. Nomenclature Practice Exercises Write the formula for the following compounds: a) iron(III) chloride b) iron(III) oxide c) lithium nitride d) magnesium nitride e) aluminum oxide Name the following compounds: a) CrCl3 b) Fe2S3 c) Ag3P d) PbS e) SnO2

24. Some Common Polyatomic Ions Formula of Ion Name of Ion OH– hydroxide ion NH4+ ammonium ion NO3–nitrate NO2– nitrite(compare to nitride: N–3) ClO3–chlorate ClO2– chlorite(compare to chloride: Cl–) CO3–2 carbonate HCO3– hydrogen carbonate or bicarbonate CN– cyanide CH3COO– acetate SO4–2 sulfate SO3–2 sulfite HSO4– hydrogen sulfate or bisulfate HSO3– hydrogen sulfite or bisulfite PO4–3 phosphate PO3–3 phosphite(compare to phosphide: P–3) HPO4–2 hydrogen phosphate H2PO4– dihydrogen phosphate A polyatomic ion is a group of atoms that as a whole has acquired a net charge. - most polyatomic ions consist of a nonmetal element such as P, S, C, or N bonded to one or more Oxygen atoms. - except for NH4+ ion, all other polyatomic ions have a negative charge of either –1, –2, or –3. - several polyatomic ions are named with ending in -ate. Related ions with one lessoxygen are named with ending in -ite. - related polyatomic ions that contain hydrogen are named by adding the word “hydrogen” before the name of the ion. If 2 hydrogens are present in the compound, the word “dihydrogen” is used.

25. Like any other anions, negatively charged polyatomic ions are associated with positive cations to balance out the charge. Ex: Na2CO3 2 Na+ + CO3–2 CaCO3 Ca2+ + CO3–2 When more than one polyatomic ion is needed to balance the charge in a given compound, the polyatomic ion is placed in parentheses with a subscript written outside to indicate the number of polyatomic ions in the compound: Ex: Mg(NO3)2 Mg2+ + 2 NO3– Al(NO3)3 Al3+ + 3 NO3– Compounds with Polyatomic Ions Ionic compounds with monoatomic ions. Ionic compounds with polyatomic ions.

26. Names of Compounds with Polyatomic Ions Compounds containing polyatomic ions are named like any other ionic compounds. When there is a metal ion with variable charge, the positive charge on the metal ion can be determined by relating it to the negative charge on the polyatomic ion: Ex: FePO4 iron(III) phosphate (Fe3+ + PO43–) Cu(NO2)2 copper(II) nitrite (Cu+ + 2 NO2–) Al2(CO3)3 aluminum carbonate (2 Al3+ + 3 CO32–) - note that Aluminum exists only in one form (Al3+), so no Roman numeral is needed. ● Write the formula for: a) Zinc phosphate b) Aluminum sulfate Aluminum sulfate is a crystalline compound used in paper making, water purification, sanitation, and tanning. ALUM (aluminum sulfate) is widely applied as inorganic coagulant for clarifying waste water treatment to make water clear for industrial and drinking. ALUM also used as mordant to allow dyes absorb by wool fibers in traditional dying. After proper biological treatment and addition of aluminum sulfate, treated waste water flows into settling basins known as clarifiers. The heavier solid particles are settled out and collected to be used as activated sludge or dewatered and land applied. The remaining clear water is then disinfected and flows to the river or ocean.

27. Covalent Compounds In ionic compounds, atoms achieve the octet rule by losing or gaining e–, in the process they acquire a net charge and the bonding between them is the attractive force between a positively charged particle and a negatively charged particle. We have seen that ionic compounds are formed between a metal and a non-metal. Atoms of 2 nonmetal elements also come together to form compounds. However, instead of gaining or losing electrons to achieve the octet rule, nonmetal atoms share one or more pair of electrons between them when forming compound, with the same end result is to acquire the noble gas configuration of 8 valence shell e–.

28. Diatomic Covalent Molecules

29. PERIODIC TABLE

30. Covalent Compounds with Single Bonds The number of covalent bonds a nonmetal atom can form is usually the number of electrons needed for the atoms to acquire a noble gas configuration.

31. Covalent Compounds with Multiple Covalent Bonds Multiple covalent bonds also form when atoms share 2 or 3 pairs of electrons to achieve their octet configuration. - Atoms of C, O, N, andS are most likely to form multiple bonds. - whereas ionic compoundsare generically referred to assalts; by contrast, covalent compoundsare calledmolecules.

32. (1) the 1st nonmetal element is named by its elemental name, and the 2nd nonmetal element is named with an ending in -ide. (2) a subscript is used to indicate two or more atoms of a given element. (3) prefixes are used to enumerate the number of atoms of a given element as indicated by its subscript. This is needed because several different compounds can be formed from the same two non-metal elements. Examples: Naming Covalent Compounds Formula Name CO Carbon monoxide CO2 Carbon dioxide SiF4 Silicon tetrafluoride N2O3 Dinitrogen trioxide P2O5 Diphosphorus pentoxide SF6 Sulfur hexaflouride P4O6 Tetraphosphorus hexoxide Note that in naming ionic compounds, one does not use prefixes to enumerate the number of different types of atoms or group of atoms in the formula. This must be deduced from the charge balance (for example, sodium carbonate = Na2CO3) - by contrast, in naming covalent compounds, one must use prefixes to indicate the atoms of the same element in a formula.

33. When a covalent bond is formed between 2 atoms of the same element, the e– pair in the bond is shared equally between the 2 atoms. Ex: In molecular hydrogen (H2), the bonding e– pair is shared equally between the two Hydrogen atoms. The e–’s spend the same amount of time in the vicinity of each atom (i.e. one can say that the e– pair is located midway between the 2 Hydrogen nuclei). A bond formed by an equally shared pair of electrons is said to be non-polar. Non-Polar Covalent Bonds - likewise, in molecular fluorine (F2), the bonding e– pair is shared equally between the 2 fluorine atoms. As a result, the F–F bondis non-polar:

34. When the 2 bonding atoms are of different elements, the bonding e– pair may not be shared equally. Ex: In the hydrogen fluoride (HF) molecule, the bonding e– pair is drawn closer to the F nucleus than to the H nucleus. This is because the F nucleus has stronger attraction to orbiting e–’s than the H nucleus. Generally within a given Period, the nuclear attraction for electrons increases in atoms with increasingly higher atomic number. We have seen that the atomic radius decreases going across a row from left to right because of increasing nuclear attraction for the orbiting electrons. - as the bonding e– pair is drawn closer to the F nucleus, they cause the F atom to develop a partial (–)charge and the H atom to develop a partial (+)charge. The character of a partial charge is often indicated by the Greek sign. Polar Covalent Bonds  The development of partial charges causes the covalent bond to become polar. In a polar covalent bond, there is a partial separation of (+) and (–) charges at either ends. Because the bond develops 2 poles, it is said to have a dipole, which is often indicates by an arrow pointing from the (+) end to the (–) end: + H –– F –

35. The propensity of an atom to draw electrons toward itself --- that is, its affinity for e– --- is a property of atoms called electronegativity. The higher the electronegativity of an atom, the higher is its affinity for electrons. - an electronegativity scale was developed by Linus Pauling and is known as the Pauling electronegativity scale, which has values ranging between 0.0  4.0. Electronegativity Linus Pauling 1901-1994

36.  The electronegativity value increases from left to right across a period and also increases going up a group from bottom to top. Fluorineis the most electronegative element, followed by Oxygen; next are Nitrogen and Chlorine. The metal elements have the lowest electronegativity values, with Cesium and Francium at the lower left corner of the periodic table having the lowest values. The electronegativity of the transition metals are generally also low, but we will not be concerned with them here.

37. The electronegativity values can be used to predict whether a bond between two atoms is polar, nonpolar, or ionic. H : HH :ClNa+ Cl–nonpolar polar ionic Predicting Bond Type from Electronegativity Values Generally, if the difference in electronegativity values in a bond is a) 0.0 to 0.4:non-polar covalentEx: C−H (2.5 − 2.1 = 0.4) b) 0.5 to 1.8:polar covalentEx: O−H (3.5 − 2.1 = 1.4) c) 1.9 or larger:ionic bondEx: K−Cl (3.0 − 0.8 = 2.2)

38. Summary of Nonpolar, Polar, and Ionic Bonds

39. Valence Shell Electron Pair Repulsion (VSEPR) Theory Molecular compounds or polyatomic ions exist in space 3-dimensionally and thus have a 3-dimensional shape. The shape a particular group of connecting atoms assumes depends on twomajor factors: 1) The number of covalent bonds that the central atom within the group forms with other atoms. 2) The electronic configuration of the valence shell e–’s surrounding the central atom; bothbonding and non-bonding valence e– pairs affect the molecular geometry of the molecule or polyatomic ion. Linear Tetrahedral Trigonal planar

40. VSEPR theory is commonly used to predict the molecular geometry of simple molecules and polyatomic ions. The theory proposes that the valence e– pairs of the central atom of a molecule or polyatomic ion tends to stay as far apart as possible from one anotherin order tominimize the electrostatic repulsion between the negatively charged electrons. As they do so, they affect the geometry of the molecule or polyatomic ion as a whole. Molecular geometry: Linear When the central atom has 2 bonding e– pairs in its valence shell, these tend to be as far apart as possible from one another due to like-charge repulsion by assuming a linear arrangement that gives a 180 angle between the 2 e– pairs. Beryllium chloride

41. Molecular geometry: Trigonal planar When the central atom has 3 bonding e– pairs in its valence shell, these will arrange into a plane directed toward the 3 corners of a flat triangle, giving rise to the molecular geometry trigonal planar with the 3 e– pairs at 120 angle relative to one another. Boron trifluoride

42. Molecular geometry: Tetrahedral When the central atom has 4 bonding e– pairs in its valence shell, these will arrange into a plane directed toward the 4 corners of a tetrahedron, giving rise to the tetrahedral shape with the 4 e– pairs at 109.5 angle relative to one another. Methane (CH4)

43. Molecular and Electronic Geometries There are 2 types of geometry that we want to look at in a molecule or polyatomic ion: 1. The electronic geometry involves all e− pairs that surround the central atom, both bonding and non-bonding e− pairs. 2. The molecular geometry involves only the bonding e– pairs of the central atom. Here the shape of the molecule or polyatomic ion is determined by the actual connection between atoms. - the next 2 slides are examples of electronic and molecular geometries of some simple molecules. Electronic geometry: Tetrahedral Molecular geometry: Trigonal pyramidal Ammonia (NH3)

44. Electronic geometry: Tetrahedral Water (H2O) Molecular geometry: Bent or Angular Methane (CH4), ammonia (NH3), and water (H2O) all have a tetrahedral electronic geometry, but each assumes a different molecular geometry based on their bonding e− pairs.

45. Using VSEPR to predict the geometry of molecule 1) First, draw the central atom in the molecule or polyatomic ion with its valence electrons as an electron-dot structure. 2) Second, determine the number of bonding e– pairs and non-bonding e– pairs (lone pairs) of this central atom. 3) Finally, determine the electronic geometry and molecular geometry with respect to the central atom. Examples: A. In BeCl2, there are 2 e– pairs in the valence shell of the central atom Beryllium. - number of bonding e–pairs =2 - number of lone e– pairs = none VSEPR predicts a molecular geometryof linear shapein which the bonding e− pairs are at opposite sides of the central atom at 180 angle from one another. - the electronic geometry is also linear.

46. B. In BF3, there are 3 e– pairs in the valence shell of the central atom Boron. - number of bonding e–pairs = 3 - number of lone e– pairs = none - molecular geometry: trigonal planarshapewith bond angles at 120. - electronic geometry: trigonal planar Carbonate ion, CO32– = trigonal planar shape A double bond is treated as if it were a single bond. – –

47. C. In ammonia (NH3), there are 4 e– pairs in the valence shell of the central atom Nitrogen. - number of bonding e–pairs =3 - number of lone e– pairs = 1 - molecular geometry: trigonal pyramidal - electronic geometry: tetrahedral D. In water (H2O), there are 4 e– pairs in the valence shell of the central atom Oxygen. - number of bonding e– pairs =2 - number of lone e– pairs = 2 - molecular geometry: bent or angular - electronic geometry: tetrahedral E. In GeF4, there are 4 e– pairs in the valence shell of the central atom Germanium. - number of bonding e– pairs =4 - number of lone e– pairs = 0 - molecular geometry: tetrahedral - electronic geometry: tetrahedral Water

48. Polarity of Molecules We have seen that the bond between 2 different atoms may be polar as a result of the difference in electronegativity value between the 2 atoms. Ex: Hydrogen Fluoride (HF) is polar because Fluorine is highly electronegative relative to Hydrogen. As a result, Fluorine pulls the bonding e– pair closer toward its nucleus and causes partial charges to develop at either ends of the bond. Because of this separation of charge in the H─F bond, Hydrogen Fluoride molecule develops a net dipole (di = 2 poles) or is said to have a dipole moment. ●Bond polarity occurs due to a separation of charges and can be described for a single bond or an entire molecule. The polarity of the molecule is the sum of all of the bond polarities in the molecule. For example, formaldehyde (CH2O) is highly polar while carbon dioxide (CO2) is nonpolar. Since CO2 is a linear molecule, the dipoles cancel each other. Formaldehyde is polar, while carbon dioxide is nonpolar.

49. The polarity of a molecule occurs when the polar bonds within it do not cancel out. By contrast, when all the polar bonds cancel out, the molecule as a whole becomes non-polar (net dipole = 0). - in predicting the polarity of a molecule, only the bonds between atoms are considered; non-bonding e− pairs (i.e. lone e− pairs) do not play a role in bond polarity. In Boron Trifluoride (BF3), symmetry of the 3 polar B─F bonds allows cancellation of one another. The molecule as a whole is non-polarand has a net dipole = 0. net dipole Water (H2O) has bent molecular geometry. The dipoles of the 2 polar O─H bonds add together to produce a net dipole over the entire molecule. Water is thus polar.

50. Carbon tetrachloride (CCl4) has a tetrahedral molecular structure. The symmetry of the 4 C─Cl bonds allows cancellation of dipoles. CCl4is nonpolar. net dipole = 0 net dipole ≠ 0 Ammonia has a trigonal pyramidal molecular structure. The polarities of the 3 N─H bonds add together to produce a net dipole over the entire molecule. Ammonia is polar. polar