IB Chemistry. Chapter 4.1

Transcription

1 IB Chemistry Chapter 4.1

2 Chemical Bonds Atoms or ions that are strongly attached to one another Chemical bonds will form if potential energy decreases (becomes more stable) 2

3 Valence Electrons Valence electrons are the electrons in the outer shell that are on the same period or row. These are the electrons that are involved in bonding atoms together. Valence electrons can easily be found by finding the group number the element is in Nitrogen has 5 valence electrons (it is in group 5) Group 1 has 1, group 2 has 2, group 3 has 3, etc.

4 Ions When an atom gains or loses an electron, it is called an ion. (The valence electrons are the ones lost) Cation if positively charged or giving electrons away. Anion if negatively charged or accepting electrons.

5 Ions Metals tend to lose their valence electrons when bonding and become cations. This is due to lower ionization energies Nonmetals tend to gain electrons in their highest energy level when bonding and become anions. This is due to high electronegativity and high ionization energies The most stable atoms have 8 electrons in their highest energy level (the octet rule)

6 Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H + Li + Na + K +

7 Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+

8 Predicting Ionic Charges B 3+ Al 3+ Ga 3+ Group 3: Loses 3 electrons to form 3+ ions

9 Predicting Ionic Charges Carbon loses 4, but the metals have multiple charges Group 4: Lose 4 electrons or gain 4 electrons?

10 N 3- P 3- As 3- Predicting Ionic Charges Nitride Phosphide Arsenide Group 5: Gains 3 electrons to form 3- ions

11 O 2- S 2- Se 2- Predicting Ionic Charges Oxide Sulfide Selenide Group 6: Gains 2 electrons to form 2- ions

12 Predicting Ionic Charges F 1- Cl 1- Br 1- I 1- Fluoride Chloride Bromide Iodide Group 7: Gains 1 electron to form 1- ions

13 Predicting Ionic Charges Group 0: Stable Noble gases do not form ions!

14 Predicting Ionic Charges Transition metals: Many transition elements have more than one possible charge. Iron(II) = Fe 2+ Iron(III) = Fe 3+

15 Predicting Ionic Charges Some transition elements have only one possible charge. Zinc = Zn 2+ Silver = Ag +

16 Polyatomic ions These are groups of atoms that hang around together and have a charge. Hydroxide ion Ammonium ion

17 Here is a trick to name and know about 48 polyatomic ions. Now to name them Like the chart above (base element name) ate BO 3 CO 3 NO 3 1 more oxygen than above Per -(base element name) ate 1 less oxygen than above (base element name ite 2 less oxygen than above Hypo (base element name) - ite PO 4 O SO 4 AsO 4 SeO 4 SbO 4 TeO 4 F ClO 3 BrO 3 IO 3 Noble Gases

18 The Octet Rule Atoms tend to gain, lose, or share e- until they are surrounded by 8 valence e- and are thus energetically stable. Exceptions do occur (and will be discussed later.) 18

19 Increasing Diff. of EN Types of Bonds 1. Ionic: electrostatic attraction between oppositely charged ions (typically between a metal and a nonmetal) 2. Covalent: sharing of e- between two atoms (typically between nonmetals) molecules created 3. Metallic: sea of e- ; bonding e- are relatively free to move throughout the 3D structure Ionic Covalent Metallic 19

20 Ionic bonding Ionic bonds do not form molecules An ionic formula is an empirical formula (smallest whole number ratio of atoms) and doesn t show what the structure looks like Ionic compounds are neutral, the positive charges balance out the negative charges. Ionic bonds form a lattice structure 20

21 Ionic Bonding Results as atoms lose or gain e - to achieve 8 electrons in the highest shell The bonded state is lower in energy (and therefore more stable). Electrostatic attraction results from the opposite charges. Occurs when diff. of EN of atoms is > 1.7 (maximum is 3.3: CsF) Can lead to interesting crystal structures Ionic compounds are brittle solids with high melting points. Solids do not conduct electricity, but molten form will conduct (ions freely moving) 21

22 Naming Ionic Compounds Cation first, then anion Monatomic cation = name of the element Ca 2+ = calcium ion Monatomic anion = root + -ide Cl - = chloride CaCl 2 = calcium chloride If the anion is a polyatomic ion, just use the name on the polyatomic ion CaSO 4 = calcium sulfate

23 Naming Ionic Compounds Metals with multiple charges (transition metals) some metal forms more than one cation use Roman numeral in name PbCl 2 Pb 2+ is cation PbCl 2 = lead(ii) chloride

24 Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Ba 2+ ( NO - 3 ) 2 Ba(NO 3 ) 2 Not balanced!

25 Writing Ionic Compound Formulas Example: Iron(III) chloride 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Fe 3+ Cl - 3 FeCl 3 Not balanced!

26 Writing Ionic Compound Formulas Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Al 3+ S Al 2 S 3 Not balanced!

27 Writing Ionic Compound Formulas Example: Magnesium carbonate 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. Mg 2+ CO 3 2- MgCO 3 They are balanced!

28 Writing Ionic Compound Formulas Example: Zinc hydroxide 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Zn 2+ ( OH - ) 2 Zn(OH) 2 Not balanced!

29 Chapter

30 Covalent Bonds Covalent bonds form molecules The formula is not always empirical but shows what the molecule looks like A molecular formula shows what the molecule actually looks like Molecular formula: C 6 H 6 empirical: CH 30

31 Molecules- 2 or more atoms bound together that act as a single, distinct object Theses molecules are shown in ball and stick form. These are also represented in structural formulas like this: H O H N H H H H H C H H

32 Condensed structural formula Ball and Stick Structural Condensed

33 Diatomics, H N F O I Cl Br Elements can exist as molecules. These elements can exist as shown. Notice that P and S have been added

34 Covalent Bonding Atoms share e- to achieve noble gas configuration that is lower in energy (and therefore more stable). Polar covalent: (different elements) e- pulled closer to more EN atom and are shared unequally -Nonpolar covalent: (same elements) e- shared equally 34

35 Polarity: Polarity in covalent bonds 1. Polar bond: bonds in which electron density is unsymmetrical 2. Dipole: Contains a + and a - pole 3. Nonpolar bond: two atoms joined are identical. 4. Covalent bonds between two unlike atoms are always polar. 5. The extent of the polarity of a covalent bond is related to the difference in the electronegativity.

36 H 2 nonpolar; the hydrogens share the electrons equally HF polar: fluorine pulls the electrons closer so they share the electrons unequally In a polar molecule, one end is partially positive and one is partially negative (Dipole) s + s - + H F or H F (vector points to neg. end) 36

37 Covalent bonds a line between atoms shows that 2 electrons are being shared H F (single bond) Multiple bonds A double line shows that 4 electrons are being shared O=O (double bond) A triple line shows that 6 electrons are being shared N=N (triple bond) Bond length : triple < double < single Bond energy : triple>double>single 37

38 Bond Order An indication of bond strength and bond length Single bond: 1 pair of e- shared Ex: F 2 :F-F: Double bond: 2 pairs of e- shared Longest, weakest Ex: O 2 O=O Triple bond: 3 pairs of e- shared Ex: N 2 :N N: Shortest, strongest 38

39 Lewis Structures Lewis structures show how valence electrons are arranged among atoms in a molecule. Lewis structures Reflect the central idea that stability of a compound relates to noble gas electron arrangemnent. Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - )

40 Lewis structures Central atom least electronegative atom (carbon is always in the center if its in the molecule) Hydrogen can only have a single bond because it only needs 2 electrons (can never be a central atom) Count the number of valence electrons in the molecule and make sure that many are in the Lewis structure. If there is a charge, add e- (if an anion) or subtract e- (if a cation). Complete the octet for all atoms (remember hydrogen only needs 2 electrons) Another exception to the rule: Boron only wants 6 electrons around it

41 .... Completing a Lewis Structure -CH 3 CCl Make carbon the central atom (least electronegative atom) Add up available valence electrons: C = 4, H = (3)(1),Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons H H.. C.. H.. Cl.. Lone pairs

42 Lone pairs- electrons that are not shared but belong to a single atom in a Lewis structure If there are too many electrons when drawing a Lewis structure, try multiple bonds

43 Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures.

44 Bond Order & Resonance Structures Bond order: single bond = 1, double bond=2, triple bond = 3 To determine bond order with resonance structures: Pick one bond and add up the integer bond order in one resonance structure to the same bond position in all other resonance structures. Divide the sum by the number of resonance structures to find bond order. 44

45 Which has shorter bonds? What is the bond order in each? SO 3 or SO 3 2- Answer: SO 3 Bond order for SO 3 is 1 1/3 bond order of SO 3 2- is 1 45

46 Molecular shapes: V.S.E.P.R. model Valence-shell electron-pair repulsion theory Because e - pairs repel, molecular shape adjusts so the valence e - pairs are as far apart as possible around the central atom. Electron domains: areas of valence e - density around the central atom; result in different molecular shapes Includes bonding e - pairs and nonbonding e - pairs A single, double, or triple bond counts as one domain A lone pair counts as one domain Summary of L m AB n : L = lone or non-bonding pairs A = central atom B = bonded atoms

47 basic domains Linear bond angle = 180 Trigonal planar bond angle = 120 Tetrahedral bond angle = If there are only two atoms in a molecule, the geometry is linear

48 # of e - domains Molecular geometry formula Predicted bond angle(s) Example (Lewis structure with molecular shape) X 2 A 180º X Linear AB 2 CO 2

49 B B A B 3 X A X X AB 3 Trigonal planar 120 NO 3 - B : A B < 120º Trigonal planar LAB 2 Bent NO 2 1-

50 B B A 109.5º 4 X X A X X B B AB 4 Tetrahedral B : A < 109.5º CH 4 or X A X B LAB 3 B Trigonal pyramidal B : Ex: NH 3 = 107º <<109.5º NH 3 X X Tetrahedral : A L 2 AB 2 B Bent Ex: H 2 O = 104.5º H 2 O

51 Molecular Polarity A molecule is polar if its centers of (+) and (-) charge do not coincide. A bond s polarity is determined by the difference of EN between atoms in bond. Partial (+) and partial (-) on atoms in a polar bond can be represented as d + and d -. Bond polarity is most often represented by an arrow that points toward the d - (most EN atom), showing the shift in e - density. d + d - H-Cl: : : H-Cl: : :

52 The sum of the bond s dipoles in a molecule determines the overall polarity of the molecule. Shapes whose dipoles cancel out (nonpolar) as long as all outside atoms are the same: Linear, trigonal planar, tetrahedral- no lone pairs (if one of the outside atoms are different, the molecule is polar) Shapes whose dipoles don t cancel out (polar): Bent and trigonal pyramidal lone pairs on central atom

53 Ex: Draw Lewis structures and name the molecular geometry. Is it polar or nonpolar? CO 2 Linear(nonpolar) H 2 S bent (polar) CCl 4 SO 2 Tetrahedral (nonpolar) bent (polar)

54 Intermolecular forces Chapter 4.3

55 Intermolecular Forces Forces that attract molecules to other molecules. These include: Hydrogen bonding Dipole-dipole attraction van der waals forces

56 Melting points and boiling points reflect the strength of bonds or intermolecular forces. Intermolecular forces are much weaker than chemical bonds

57 Relative Magnitudes of Forces The types of bonding forces vary in their strength as measured by average bond energy. Strongest Covalent bonds (400 kcal) Hydrogen bonding (12-16 kcal ) Dipole-dipole interactions (2-0.5 kcal) Weakest Van der waals (less than 1 kcal)

58 Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole. H F d+ d-

59 Dipole-Dipole Attraction Attraction between oppositely charged regions of neighboring molecules. As the polarity increases, the attraction increases. Dipole-dipole occurs in polar molecules

60 Hydrogen Bonding attraction between hydrogen and more electronegative neighboring atoms such as fluorine, oxygen and nitrogen. (very polar bonds) This is just a strong dipole-dipole attraction Hydrogen bonding occurs in polar molecules where H is bonded to F, O, or N

61 Hydrogen Bonding in Water

62 Induced or instantaneous dipoles Occurs in all molecules and atoms Van der waals forces These are the only forces of attraction between completely nonpolar molecules Larger molecules (more electrons) create stronger van der waals forces

63 Van der waals in Hydrocarbons

64 List the following in order of increasing boiling points: BaCl 2, H 2, CO, HF, Ne Answer: H 2, Ne, CO, HF, BaCl 2 nonpolar, nonpolar, polar, polar, ionic Van der waals, van der waals, dipole-dipole, hydrogen, ionic

65 Chapter 4.4

66 Metallic bonding Metallic elements have low I.E.; this means valence e - are held loosely. A metallic bond forms between metal atoms because of the movement of valence e - from atom to atom to atom in a sea of electrons. The metal thus consists of cations held together by negatively-charged e - "glue. This results in excellent thermal & electrical conductivity, ductility, and malleability. A combination of 2 metals is called an alloy. 66

67 Free e - move rapidly in response to electric fields, thus metals are excellent conductors of electricity. Free e - transmit kinetic energy rapidly, thus metals are excellent conductors of heat. Layers of metal atoms are difficult to pull apart because of the movement of valence e-, so metals are durable. However, individual atoms are held loosely to other atoms, so atoms slip easily past one another, so metals are ductile.

68 Review of bonding Ionic compound (metal/nonmetal) creates a lattice Formula doesn t tell the exact number of atoms, but instead the smallest whole number ratio of atoms (no lewis structures) Covalent compound (nonmetal/nonmetal) creates molecules Formula tells you the exact number of atoms in the molecule. (lewis structures)