Chapter 8 Covalent Boding

Transcription

1 Chapter 8 Covalent Boding

2 Molecules & Molecular Compounds In nature, matter takes many forms. The noble gases exist as atoms. They are monatomic; monatomic they consist of single atoms. Hydrogen chloride (HCl) is a gas at room temperature. Water is a liquid at room temperature. Salts are crystalline solids with high melting points. Compounds such as HCl and water are not ionic. Their combining atoms do not give up electrons or accept electrons as readily as sodium does in combining with chlorine.

3 Molecules & Molecular Compounds Instead, a tug-of-war for the electrons takes place between the atoms, bonding the atoms together. The atoms held together by sharing electrons are joined by a covalent bond. bond Many elements found in nature are in the form of molecules. A molecule is a neutral group of atoms joined together by covalent bonds. Example is oxygen molecules. Each oxygen molecules consists of two oxygen atoms joined by covalent bonds. A diatomic molecule is a molecule consisting of two atoms of the same element.

4 Molecules & Molecular Compounds There are seven elements that are always found as diatomic molecules Iodine, Hydrogen, Nitrogen, Bromine, Oxygen, Chlorine, Fluorine I Have No Bright Or Clever Friends Another device is your friend's name, Cl I F H Br O N

5 Molecules & Molecular Compounds Atoms of different elements can combine chemically to form compounds. In many compounds, atoms are bonded to each other to form molecules. (H2O, CO) A compound composed of molecules is called a molecular compound. The molecules of a given molecular compound are all the same. There is not such thing as a molecule of sodium chloride (NaCl) or magnesium chloride. (MgCl2) These ionic compounds exist as collections of + and charged ions arranges in repeating three dimensional patterns.

6 Molecules & Molecular Compounds Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds. Many molecular compounds are gases or liquids at room temperature. Ionic compounds are formed from a metal combined with a nonmetal. Most molecular compounds are composed of atoms of two or more nonmetals.

7 Molecular Formulas A molecular formula is the chemical formula of a molecular compound. A molecular formulas shows how many atoms of each element of a molecule contains. A water molecule consists of two hydrogen atoms and one oxygen atom. The molecular formula of water is H2O. The subscript written after the symbol indicates the number of atoms of each element in the molecule. (If there is only 1 atom, the subscript 1 is omitted.) The molecular formula of carbon dioxide is CO2.

8 Molecular Formulas Ethane is C2H6. Ethane contains 2 carbon atoms and 6 hydrogen atoms. A molecular formula reflects the actual number of atoms in each molecule. The subscripts are not necessarily lowest whole-number ratios, as in ionic formula units. Molecular formulas also describe molecules consisting of one element. Example O2 A molecular formula does not tell you about a molecule s structure. It does not show either the arrangement of the various atoms in space or which atoms are covalently bonded to one another.

9 Questions How are the melting points and boiling points of molecular compounds usually different from ionic compounds? Molecular compounds tend to have relatively lower melting points and boiling points. What information does a molecular formula provide? Show how many atoms of each element one molecules of a compound contains. What are the only elements that exist in nature as uncombined atoms & what term is used to describe them? Noble gases, monatomic.

10 End of Section 8.1

11 Octet Rule for Covalent Bonding Recall that in forming ionic compounds, electrons tend to be transferred so that each ion acquires a noble gas configuration. In forming covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. Example: each H atom has one electron. But a pair of H atoms share these two electrons when they form a covalent bond in the H molecule. Each H atom thus attains the electron configuration of He, a noble gas with two electrons.

12 Covalent Bonding Combinations of atoms of the nonmetallic elements in Groups 4A, 5A, 6A, and 7A are likely to form covalent bonds. The atoms usually acquire a total of 8 electrons, or an octet, by sharing electrons so that the octet rule applies. The hydrogen atoms in a hydrogen molecule are held together mainly by the attraction of the shared electrons to the positive nuclei. Two atoms held together by sharing a pair of electrons are joined by a single covalent bond.

13 Covalent Bonding The halogens (7A) also form single covalent bonds in their diatomic molecules. Chlorine atom has 7 valence electrons (group 7A), it needs one more to attain an octet. By sharing electrons and forming a single covalent bond, two chlorine atoms achieve an octet.

14 Covalent Bonding A pair of valence electrons that is not shared between atoms is called an unshared pair. (or lone pair or nonbonding pair)

15 Covalent Bonding Take a look at ammonia (NH3). The ammonia molecules has one unshared pair of electrons.

16 Covalent Bonding Another example, methane (CH4)

17 Practice Problems Draw electron dot structures for each molecule: chlorine, bromine, iodine Draw an electron dot structure for each of the following molecules that have single covalent bonds: H2O2, PCl3

18 Double and Triple Covalent Bonds Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons. Double Covalent Bond a bond that involves two shared pairs of electrons Triple Covalent Bond a bond formed by sharing three pairs of electrons.

19 Double and Triple Covalent Bonds Single, double and triple covalent bonds can also exist between unlike atoms. CO2-2 oxygen atoms (6 valence e-), each share 2 e- with carbon (4 valence e-) to form a total of two carbon-oxygen double bonds. Carbon dioxide is an example of a triatomic molecule, molecule which is a molecule consisting of three atoms. They do not have to be the same Elements.

20 Coordinate Covalent Bonds Carbon monoxide (CO) is an example of a type of covalent bonding different from that seen in water, ammonia, methane and carbon dioxide. A carbon atoms needs to gain four electrons to attain an octet. An oxygen atom needs two electrons. Coordinate covalent bond is a covalent bond in which one atom contributes both bonding electrons. CO has 2 covalent bonds and 1 coordinate covalent bond

21 Coordinate Covalent Bonds If the structure is a molecular ion, add one valence electron for each negative charge and remove one valence electron for each positive charge. Polyatomic ion is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit. (NH4+, SO42-, PO43-)

22 Polyatomic Ions Most polyatomic cations and anions contain both covalent and coordinate covalent bonds. Therefore, compounds containing polyatomic ions include both ionic and covalent bonding. The electron dot structure for a neutral molecule contains the same number of electrons as the total number of valence electrons in the combining atoms. The negative charge of a polyatomic ions shows the number of electrons in addition to the valence electrons of the atoms present. Because a negatively charged polyatomic ion is part of an ionic compound, the positive charge of the cation of the compound balances these additional electrons.

23 Practice Draw the electron dot structure of the following: Hydroxide ion (OH-) The polyatomic boron tetrafluoride anion (BF4-) Sulfate - SO42- (S is the central atom) Carbonate - CO32- (C is the central atom) Hydrogen carbonate ion - HCO3attached to O) (C is central & H is

24 Bond Dissociation Energies A large quantity of heat is given off when hydrogen atoms combine to form hydrogen molecules. This suggests that the product is more stable than the reactants. The covalent bond in the hydrogen molecule (H2) is so strong that it would take 435 kj of energy to break apart all of the bonds in 1 mole (a very large number) of H2. The energy required to break the bond between two covalently bonded atoms is the bond dissociation energy. Usually expressed as the energy needed to break one mole of bonds. (kj/mol)

25 Bond Dissociation Energies A large bond dissociation energy corresponds to a strong covalent bond. A carbon-carbon single bond has a bond dissociation energy of about 347 kj/mol. Strong carbon-carbon bonds help explain the stability of carbon compounds. Bond dissociation energy is a measure of bond strength Bond strength - Triple > double > single Bond dissociation energy Triple > double > single

26

27 Resonance Ozone in the upper atmosphere blocks harmful ultraviolet radiation from the sun. At the lower elevations, it contributes to smog. Ozone molecule has two possible electron dot structures. The electron pairs were thought to rapidly flip back and forth, or resonate, resonate between the different electron dot structures. Ozone consists of 1 single coordinate covalent bond, and 1 double covalent bond.

28 Resonance Double covalent bonds are usually shorter than single bonds, so it was once thought that the bond lengths in ozone were unequal. Experiments showed that was not the case, however. The two bonds are the same length. The actual bonding in the ozone molecule is the average of the two electron dot structures. The electron pairs do not actually resonate back and forth. The actual bonding of oxygen atoms in ozone is a hybrid, or mixture of the extremes represented by the resonance forms.

29 Resonance Resonance structure is a structure that occurs when it is possible to write two or more valid electron dot formulas that have the same number of electron pairs for a molecule or ion. Resonance structures are a way to envision the bonding in certain molecules. Although no back-and-forth changes ocurr, doubleheaded arrows are used to connect resonance structures.

30

31

32

33

34

35

36

37 Questions What electron configurations do atoms usually achieve by sharing electrons to form covalent bonds? Noble gases configurations How is an electron dot structure used to represent a covalent bond? Two dots represent each covalent bond When are two atoms likely to form a double bond between them? A triple bond? When they can attain a noble gas structure by sharing two pairs or three pairs of electrons.

38 Questions How is a coordinate covalent bond different from other covalent bonds? The shared electron pair comes from one of the bonding atoms. In other covalent bonds, each bonding atom provides an electron. How is the strength of a covalent bond related to its bond dissociation energy? A large bond dissociation energy corresponds to a strong covalent bond.

39 Questions List three ways in which the octet rule can sometimes fail to be obeyed. Molecules whose total number of valence electrons is an odd number. Atom has fewer or more than a complete octet of valence electrons. Draw electron dot structures for the following molecules, which have only single covalent bonds. H2S, PH3, CIF Which bond is stronger? H2 has a dissociation energy of 435 kj/mol, a carbon-carbon bond has a dissociation energy of 347 kj/mol The H-H bond is stronger because it has a greater dissociation energy.

40 End of Section 8.2

41 Molecular Orbitals The model for covalent bonding we have been using assumes that the orbitals are those of the individual atoms. There is a quantum mechanical model of bonding, however, that describes the electrons in molecules using orbitals that exist only for groupings of atoms. When two atoms combine, the model assumes that their atomic orbitals overlap to produce molecular orbitals that apply to the entire molecule. In some ways, atomic orbitals and molecular orbitals are similar.

42 Molecular Orbitals Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole. Each atomic orbital is filled if it contains two electrons. Similarly, two electrons are required to fill a molecular orbital. A molecular orbital that can be occupied by two electrons of a covalent bond is called a bonding orbital. orbital

43 Molecular Orbital - Sigma Bond When two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei, a sigma (σ) bond is formed. The distinguishing feature of a sigma bond is that the overlap region lies directly between the two nuclei. It does not matter what shapes the orbitals have or what types they are. They can be s orbitals or p orbitals or hybrid orbitals.

44 Molecular Orbitals The side by side overlap of atomic p orbitals produces pi (π) molecular orbital. When a pi molecular orbital is filled with 2 electrons, a pi bond results. The pi bond has orbital overlap off to the sides of the line joining the two nuclei.

45 Molecular Orbitals Atomic orbitals in pi bonding overlap less than in sigma bonding. Therefore, pi bonds tend to be weaker than sigma bonds.

46 VSEPR Theory Electron dot structures fail to reflect the three dimensional shapes of the molecules. Valence Shell Electron Pair Repulsion Theory (VSEPR) According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence electron pairs stay as far apart as possible. Unshared pairs of electrons no bonding atom is vying for these unshared electrons so they are held closer to the central atom than are the bonding pairs. The unshared pair strongly repels the bonding pairs pushing them together.

47 VSEPR Theory

48

49 VSEPR Theory Steric number = (number of lone pairs on central atom) + (number of atoms bonded to central atom) The steric number is determined from the Lewis structure. Steric number determines the electron-pair arrangement, the geometry that maximizes the distances between the valence-shell electron pairs.

50

51 VSEPR In predicting molecular shapes, it may be useful to start with an electron dot structure. The electron dot structure shows both the bonding and nonbonding pairs of electrons around the central atom. When using VSEPR theory to predict molecular shape, double and triple bonds are viewed as single bonds. When 4 pairs of electrons must be accommodated around the central atom tetrahedral (109.5º) When 3 pairs a trigonal planar maximizes space (120º) When 2 pairs a linear arrangement (180º)

52 Hybrid Orbitals VSEPR works well when accounting for molecular shape, but does not help much in describing the types of bonds formed. Orbital hybridization provides information about both molecular bonding and molecular shape. In hybridization, hybridization several atomic orbitals mix to form the same total number of equivalent hybrid orbitals.

53

54 End of Section 8.3 End of Chapter 7

55 Bond Polarity Covalent bonds involve electron sharing between atoms. However, they differ in terms of how the bonded atoms share the electrons. The bonding pairs of electrons are pulled between the nuclei of the atoms sharing the ewhen the atoms in the bond pull Equally (identical atoms are bonded), the bonding electrons are shared equally and the bond is a nonpolar covalent bond. H2, O2, N2, and Cl2 all have nopolar covalent bonds

56 Bond Polarity A polar covalent bond (polar bond), is a covalent bond between atoms in which the electrons are shared unequally. The more electronegative atoms attracts electrons more strongly and gains a lightly negative charge. The less electronegative atoms has a slightly positive charge. Fluorine is the most electronegative and oxygen is the second most electronegative element. Electronegativity describes the attraction an atom has for electrons when the atom is in a compound.

57

58 Electronegativity Values of Some Elements

59 Bond Polarity In hydrogen chloride (HCl) molecule, the H has an electronegativity of 2.1 and the chlorine has an electronegativity of 3.0. The values are significantly different, so the covalent bond in HCl is polar. The chlorine atom acquires a slightly negative chare and the hydrogen atoms acquires a slightly positive charge. Greek letter delta (δ) indicated that the atoms involved in the covalent bond acquire only partial charges, less than 1+ or 1-. δ+ δ- H Cl δ+ δ- δ- H O H δ+ δ- H - Br

60 Bond Polarity The electronegativity difference between two atoms tells you what kind of bond is likely to form. As the electronegativity difference between two atoms increases, the polarity of the bond increases. If the electronegativity difference is greater than 2.0, it is very likely that electrons will be pulled away completely by one of the atoms. (an ionic bond will form) EN Differences Probable Bond Type Example Nonpolar covalent H H (0.0) Moderately polar covalent H Cl (0.9) Very polar covalent H F (1.9) > 2.0 ionic Na Cl (2.1)

61 Problems Place the following covalent bonds in order from least to most polar. a. H Cl b. H Br c. H S d. H - C c & d (tie), b, a Identify the bonds between atoms of each pair of elements as nonpolar covalent, moderately polar covalent, very covalent, or ionic. a. H & Br e. Li & O b. K & Cl f. Br & Br c. C & O d. Cl & F a. moderately polar covalent b. ionic c. moderately to very polar covalent d. moderately to very polar covalent e. ionic

62 Polar Molecules The presence of a polar bond in a molecule often makes the entire molecule polar. In a polar molecule, molecule one end of the molecule is slightly negative and the other end is slightly positive. In HCl, the partial charges on the hydrogen and chlorine atoms are electrically charged regions or poles. A molecule that has two poles is called a dipolar molecule. There is no monopolar or tripolar molecule. When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates.

63 Polar Molecules The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds. CO2 has two polar bonds and is linear. The C and O lies along the same axis, thus the bond polarities cancel because they are in opposite directions. CO2 is a nonpolar molecule despite the presence of two polar bonds. Water also has two polar bonds. The water molecule is bent and the bond polarities do not cancel and the water molecule is polar.

64 Attractions Between Molecules Molecules can attract each other by a variety of forces. Intermolecular attractions are weaker than either ionic or covalent bonds. Intermolecular forces are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature. Van der Waals Forces are the two weakest attractions between molecules and include both: dipole interactions dispersion forces.

65 Dipole Interactions Dipole interactions occur when polar molecules are attracted to one another. The attraction involved occurs between the oppositely charged regions of polar molecules. The slightly negative region is weakly attracted to the slightly positive regions of another molecule. Dipole interactions are similar to but much weaker than ionic bonds.

66 Dipole Interactions

67 Dispersions Forces Dispersion forces are the weakest of all molecular interactions and are caused by the motion of electrons. They occur between nonpolar molecules. When the moving electrons happen to be momentarily more on the side of a molecule closest to a neighboring molecule, their electric force influences the neighboring molecule s electrons to be momentarily more on the opposite side. This causes an attraction between the two molecules similar but much weaker than the attraction between polar molecules.

68 Dispersions Forces The strength of dispersion forces generally increases as the number of electrons in a molecule increases. The halogen diatomic molecules attract each other mainly by means of dispersion forces. Fluorine and chlorine has relatively few electrons and are gases at room temperature because of their weak dispersion forces. Bromine has a larger number of electrons and generates larger dispersion forces. Bromine molecules attract each other sufficiently to make it a liquid at room temperature. Iodine, with a larger number of electrons, is a solid at room temperature.

69 Dispersions Forces

70 Hydrogen Bonds The dipole interactions in water produce an attraction between water molecules. Each O H bond in the water molecule is highly polar and the O acquires a slightly negative charge. The positive region of one water molecules attracts the negative region of another water molecule. This attraction between the hydrogen of one water molecule and the oxygen of another water molecule is strong compared to other dipole interaction. This strong attraction, (also found in other H containing molecules) is called hydrogen bonding.

71 Hydrogen Bonds

72 Hydrogen Bonds Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom. The other atom may be in the same molecule or in a nearby molecule. Hydrogen bonding always involves hydrogen. The bond between hydrogen and a very electronegative atom is strongly polar. Hydrogen bonds are the strongest of the intermolecular forces.

73 End of Chapter 8