Unit 4: Chemical Bonds. Chapter 7-9

Transcription

1 Unit 4: Chemical Bonds Chapter 7-9

2 Objectives 26 Identify the number of valence electrons for elements and their Lewis dot structure 27 Define the terms cation and anion including radius size and charge 28 Determine the isoelectronic electron configurations for atoms and their ions including the ionic charges 29 Identify the properties of ionic bonds 30 Predict the shape of molecules using the VSEPR theory 31 Identify the bonds between certain elements within a compound as non-polar, polar, or ionic 32 State and identify the three intermolecular forces including London dispersion forces and how they affect melting points, dipole forces, and hydrogen bond forces 33 Convert between formula and chemical name for covalently bonded molecules, binary ionic compounds, polyatomic ionic compounds, and hydrates 34 Identify the dissociation factor of compounds

3 26 Valence Shells Before we discuss bonds, we need to determine the number of electrons in the valence shell of an ion. This is accomplished by looking at the electron configuration.

4 Valence Shells O: 1s 2 2s 2 2p 4 Consider oxygen It s outer energy level is the second energy level. This tells us there are 6 valance electrons.

5 Lewis Dot Structures Once we know the number of valence electrons, it is possible to give a visual representation. This is called a Lewis Dot structure. It requires a dot for each valence electron surrounding the element symbol. Paired electrons should still be depicted.

6 Lewis Dot Structures Oxygen has 6 valence electrons. If we look at the orbital notation for just its valence shell, we get the following: 2p 2s Its Lewis Dot structure would look as follows. O

7 Lewis Dot Structures O Notice how oxygen has two electrons that are not paired up. This indicates that oxygen would like to gain two more electrons so it has 8 total electrons. All elements are most stable with 8 valence electrons. This is known as the Octet Rule. There are 5 exceptions: hydrogen, helium, lithium, beryllium, and boron. These prefer to have 2 electrons in their valence shell.

8 27 Cations and Anions Atoms can gain or lose electrons. When atoms lose electrons, they become positive and are called cations. When atoms gain electrons, they become negative and are called anions.

9 Ionic Radii Nuclear charge holds electrons a certain distance from the nucleus. As a cation is formed, there are less electrons for the nucleus to hold. This allows the nucleus to pull the outer energy levels slightly closer.

10 Ionic Radii As an anion is formed, there are more electrons for the nucleus to hold. The nucleus does not have enough charge to hold the extra electrons as close, and as a result, the radius increases slightly.

11 28 Isoelectronic Configurations As discussed in Unit 3, each atom has an electron configuration to show where each electron belongs. For example: Al: 1s 2 2s 2 2p 6 3s 2 3p 1 When ions are formed, the electrons are either added to the last energy levels or taken from the last energy levels. For example: assume we take three electrons from aluminum. Al +3 : 1s 2 2s 2 2p 6 Aluminum s 3s and 3p orbitals are now empty.

12 Isoelectronic Configuration The term isoelectronic can be broken down into: Iso: same Electronic: electrons Therefore, the term means the same electron configuration as another element. In chemistry, this refers to a noble gas.

13 Isoelectronic Configuration If we look at an element, the number of electrons it holds are close to a noble gas. This means it will tend to gain or lose electrons until it matches that noble gas. For example: Oxygen has 8 electrons and is thus close to neon s 10. Neon has a configuration of 1s 2 2s 2 2p 6. Oxygen has a configuration of 1s 2 2s 2 2p 4. For oxygen to be isoelectronic with neon, it needs two more electrons, thus oxygen tends to gain electrons to get: O -2 : 1s 2 2s 2 2p 6

14 Charges The elements marked below will always carry the charge indicated. The elements in white can have charges that vary. These will be determined with a Roman numeral. Transition Metals

15 29 Ionic Properties Every ionic compound will follow certain properties. They are: Ionic compounds form crystalline structures. Ionic compounds are brittle. Ionic compounds have high melting and boiling points. Ionic compounds as solids will not conduct electricity. Ionic compounds in solution will conduct electricity.

16 30 Molecular Shapes Covalently bonded molecules will create different shapes. These shapes are controlled by the bonds formed and the paired electrons. To determine the shapes, the valence shell electron pair repulsion (VSEPR) theory is used.

17 VSEPR Theory VSEPR Theory states that the valence electrons in a molecule will position themselves so they are as far away from the other electrons as possible. To determine the shapes of molecules using this theory, the number of bonds and pairs of electrons must be determined.

18 VSEPR Theory Recall Lewis Dot structures from the previous unit. Each dot represents a valence electron. Oxygen has six valence electrons so its Lewis Dot structure is as follows: O These paired electrons are called lone pairs because they belong to only one atom.

19 VSEPR Theory The Lewis Dot structures can be used to determine the bonds created. Take water which is H 2 O. O H O H H H When a bond is formed, the electrons are shown with a single line between the atoms.

20 VSEPR Theory Notice that the water molecule still contains two sets of lone pair electrons. These electrons will force the hydrogens to create a bent shape. O H H

21 VSEPR Theory There are five shapes the basic covalent molecules will create. Each can be determined by looking at the central atom of the molecule. The two components to look at are the number of atoms bound and the number of lone pairs on the central atom.

22 VSEPR Theory Shape Lone Pairs of electrons on Central Atom Atoms bonded to the Central Atom Example Linear 0 2 CO 2 Bent 1 or 2 2 H 2 O or HNO Trigonal Planar 0 3 BF 3 Trigonal Pyramidal 1 3 NH 3 Tetrahedral 0 4 CH 4 In addition, any molecule that has only two atoms will be linear. i.e.: oxygen gas, O 2

23 31 Intramolecular Bonds When a bond is created, it is often stated that the atoms share their electrons. While this is true to some degree, the sharing is not always equal. Each atom has its own electronegativity. The tendency of an atom to attract a bonded electron to itself. The greater an atom s electronegativity, the more time the electrons will spend near that atom.

24 Bond Strength This pulling of the electrons towards one atom can create partial charges. Each bond can be classified as either nonpolar covalent, polar covalent, or ionic. To determine how atoms share, compare their electronegativity values. The greater the difference, the more ionic character will be present in the molecule.

25 Intramolecular Bonds Nonpolar Polar Covalent Ionic Covalent Difference in Electronegativity As we look at the difference in electronegativity, we can use the chart above to determine the type of bond. Take water for example: H 2 O The bonds formed are between hydrogen and oxygen. They have the following electronegativities: 2.2 for H and 3.4 for O The difference between the two is 1.2 and thus the bond is polar covalent. Because oxygen has the larger electronegativity, the electrons spend more time near oxygen. This creates a partial negative charge on the oxygen atom and a partial positive charge around the hydrogen atom.

26 32 Intermolecular Forces While intramolecular forces occur inside a molecule, intermolecular forces occur between molecules. Three intermolecular forces exist: London Forces Dipole Forces Hydrogen bonding

27 Intermolecular Force Strength Each of the intermolecular forces hold molecules together. However, certain forces are stronger. London is the weakest. Dipoles use the partial charges as an attractive force making them stronger than London. H-Bonding is the strongest because of the partial charges and the use of hydrogen.

28 Determining Intermolecular Forces To determine the intermolecular force on a molecule, it is necessary to know whether it is polar or not. If polar, the molecule will have a partial positive and partial negative side. This can be determined using the electronegativities and the Lewis Dot structure.

29 Determining Intermolecular Forces Water has a Lewis Dot structure as shown. Oxygen has an electronegativity of 3.4 while hydrogen has an electronegativity of 2.2. This would mean that oxygen is partial negative and the hydrogens are each partial positive. Because the molecule has a postive and negative side, it is considered polar. If the molecule would have had the same charge on the entire outside, it would be considered non polar.

30 Intermolecular Forces Force Strength Polar/Nonpolar Unique Characterestics H-Bonding Strongest Polar Must contain H and either O, N, F Dipole Forces Medium Polar London Forces Weak Nonpolar Dipole forces and H-Bonding are the only forces that are polar, but H-bonding has element requirements. If a molecule is polar but does not contain one of the elements listed above, it must be a dipole.

31 33 Writing Binary Formulas and Names Binary compounds refer to compounds that contain 2 elements. When writing the name of a binary compound, list the first element exactly as it appears on the Periodic Table. For the second element, drop the ending of the element s name and add ide.

32 Binary Names-Examples NaCl sodium chloride (chlorine drops the ine) CaBr 2 calcium bromide (bromine drops the ine)

33 Writing Binary Formulas and Names Writing the formulas from the names requires the use of the charges. It is important to balance the positive charge with the negative charge. This is done by adding subscripts to the elements.

34 Binary Formulas - Examples calcium phosphide Ca 3 P 2 Calcium has a +2 charge. Phosphorus has a -3 charge To balance their charges, we need to have multiple atoms. If we add another calcium, we will have an overall charge of +4 to -3. Let s add another phosphorus to give us a overall charge of +4 to -6. Since we are only off by 2, adding another calcium will balance the charges at +6 to -6.

35 Writing Formulas We got to the answer on the last slide by using logical method of adding one atom at a time. This can also be done by looking for the least common multiple. Since we had charges of 2 and 3, the least common multiple is 6. (2 x 3 = 6) (3 x 2 = 6) Therefore, the atom with the charge of 2 requires 3 atoms and the atom with the charge of 3 requires 2. Ca 3 P 2

36 Writing Formulas magnesium bromide MgBr 2 Magnesium has a +2 charge. Bromine has a -1 charge. The least common multiple is two. 1 x 2 = 2 2 x 1 = 2

37 Transition Metals If you recall from the slide on charges, transition metals did not have a defined charge. Their charges vary and thus a Roman numeral is used to determine their charge. This Roman numeral is always listed directly after the metal.

38 Transition Metals Fe 2 O 3 For the formula above, we know Fe is iron and O is oxide (oxygen as an ion) Oxygen has a -2 charge and since there are 3, this compound has a overall -6 charge. Since we have to have a +6 charge as well, we have to consider a number times 2 to give 6. (? x 2 =6) In this case, the answer would be three. Therefore, the name of the compound is iron (iii) oxide.

39 Transition Metals Manganese (ii) chloride In this compound, Manganese has a +2 charge and chlorine has a -1 charge. Therefore, the least common multiple is 2. MnCl 2

40 Polyatomic Ions Some elements will combine covalently (Unit 6) and still have a charge. As long as they have a charge, they can create ionic compounds. They are treated as though they are single entities.

41 Naming the Polyatomics From a formula, naming the polyatomic ionic compounds requires element to be named and the polyatomic ion. For instance: CaSO 4 Ca represents calcium SO 4 represents sulfate The name of this compound is calcium sulfate.

42 Writing formulas Writing formulas from the name works the same as binary compounds. Determine the charge on each. Find the least common multiple. Add the proper subscripts. If more than one polyatomic ion is required, add parenthesis around the ion before adding the subscript.

43 Writing formulas-example Iron (iii) nitrate Iron (iii) refers to Fe +3 Nitrate refers to NO 3-1 Therefore, the least common multiple is 3 and three nitrates are required. Fe(NO 3 ) 3 The parenthesis tells us that there are 3 N and 9 O.

44 Hydrates Hydrates are unique ionic compounds that attract water. Each hydrate is surrounded by a certain number of water molecules. These water molecules need to be identified in both the formula and the name.

45 Hydrates To indicate a hydrate, a dot is used to indicate a weak bond. The number of water molecules are indicated with a numeric prefix. The ionic part of the compound is named as previously described. For example: copper (ii) sulfate pentahydrate CuSO 4 5H 2 O

46 Molecular Nomenclature When naming covalent molecules, first identify each element. If there is more than one of the first element, add the appropriate prefix to the front of its name. The second element should always includes its prefix and its ending should change to ide.

47 Molecular Nomenclature carbon CO 2 di oxide To name this compound, first identify each element. Since there is only one of the first element, no prefix is needed. There are two of the second element so the prefix di- will be added. Notice the ending of oxygen has already be changed.

48 Molecular Nomenclature Writing the formulas will simply work in the opposite direction. Identify and write the symbol for each element. Use the prefixes to determine the subscript of each element.

49 Molecular Nomenclature Tetraphosphorus decoxide P 4 O 10 First, record the symbols for each element. Tetra- indicates four so the subscript on phosphorus will be four. Deca- means ten so the subscript on oxide will be ten

50 34 Dissociation Factors Dissociation factors describe how many pieces an ionic compound can divide into. This is calculated by adding the subscripts of each ion. Be careful because the one s are omitted when writing formulas.

51 Dissociation Factors Assume we have calcium chloride: CaCl 2 If this molecule breaks apart, we will have 1 calcium ion and 2 chloride ions. This means the dissociation factor is 3.

52 Dissociation Factors The same idea applies to polyatomics: Ca(NO 3 ) 2 If this molecule breaks apart, we will have 1 calcium ion and 2 nitrate ions. This means the dissociation factor is 3. The polyatomic ions do not break apart.

53 This concludes the tutorial on measurements. To try some practice problems, click here. To return to the objective page, click here. To exit the tutorial, hit escape.

54 Prefixes Return mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7 octa 8 nona 9 deca 10

55 Definitions-Select the word to return to the tutorial Valence Shell: Outer most energy level of an atom