Chemistry 20 Final Review

1. Chemistry 20 Final Review • Bonding Unit • Gases Unit • Solutions, Acids and Bases Unit • Stoichiometry Unit

2. Bonding Unit Topics: • Lewis Structures: • These are your electron diagrams for individual elements (showing valence e-) • Ex. S Mg

3. Intramolecular Forces: • Remember these are the forces WITHIN a molecule • What forces hold a molecule together? • ANS: BONDS • What type of bonds are there? • ANS: COVALENT AND IONIC

4. Ionic Bond Structures • Remember how to draw electron dot diagrams for ionic compounds • They DO NOT SHARE electrons – the metal looses its outer shell electrons and the non-metal gains to a full 8 e- • Ex:

5. Covalent Bond Structures • These are the bonds holding MOLECULAR compounds together • They DO SHARE the electrons • These were also called Lewis Structures • The element that goes in the middle is the one with the most BONDING e- • Examples:

6. H H P H H H P H eg) PH3

7. Remember double and triple bonds: • Each element except hydrogen needs 8 electrons around it and there should be NO LONE PAIRS • This is when double and triple bonds form • Ex. O O N N

8. H H C C H H Structural Diagrams and Shape Diagrams • When there are 2 shared electrons between two elements in a molecule draw a line to show this bond • Ex. • Remember the shapes and shape codes • Ex. tetrahedral, trigonal planar, pyramidal, linear, bent

9. N H H H   H C H H H • the code has two numbers: 1. the number ofattached to the central atom atoms 2. the number of on the central atom lone pairs CH4 eg) NH3(g) 3 - 1 4 - 0 pyramidal tetrahedral

10. tetrahedral 4 – 0 CH4 3 – 0 trigonal planar CH2O 3 – 1 pyramidal NH3 2 – 1 bent HNO 2 – 2 bent H2O ***all other codes are linear • Molecules take on these shapes due to the VSEPR theory - valence shell electron pair repulsion • molecules adjust their shapes so that valence e- are as far away from each other as possible

11. Polar vs. Nonpolar • Remember electronegativities: • The number in each element box above the element • It shows how badly an elements wants e- • The higher the number, the stronger it pulls • When two elements are bonded together and there is a difference in electroneg. then you have a polar bond • Ex. H – F (see next slide)

12. “arrow” points towards element with higher electronegativity (-) Bond Dipole Arrows “+” at the end that is + H – F + -

13. you can use the difference in electronegativity between two atoms to determine the bond Difference in Electronegativity 3.3 1.7 0.5 0 slightly polar covalent mostly ionic polar covalent non-polar covalent

14. Polar vs. Nonpolar Molecules

15. tetrahedral: if all atoms attached have the same pull (in or out), if different atoms attached nonpolar polar • trigonal planar: if all atoms attached have the same pull (in or out), if different atoms attached nonpolar polar • pyramidal: as long as there is a difference in electronegativity between the atoms polar • bent: polar • linear: …look at electronegativity difference polar or nonpolar

16. O H C C H H H H Cl H C C I Examples 1. H2O 2. HCl polar polar 3. C2H2 4. C2HI np np polar nonpolar

17. Intermolecular Forces • These are the forces that cause attraction BETWEEN molecules • They are weaker then bonding within a molecule • They are responsible for the bp and mp of compounds since when you boil/ melt a molecule you are ONLY breaking these forces BETWEEN molecules • The three intermolecular forces we talked about the occur between MOLECULAR compounds • HB, DD, LD

18. - + - + + - - + - + DD - Dipole - Dipole • These attractions occur in POLAR molecular compouds ONLY • The slightly positive end of one molecule is attracted to the slightly negative end of another molecule

19. LD: London Dispersion Forces • These attractive force occurs between ALLmolecular compounds • It is caused by electrons in atoms and molecules constantly being in motion • So sometimes one side of a molecule can have more electron then the other side • This creates a temporary polar molecule • An attraction then forms between the ends of these polar molecules • Remember you have stronger LD forces as the molecule becomes larger or has more electrons

20. HB: Hydrogen Bonding • These attractive forces occur in molecular compounds that H bonded to either N, O or F • Draw the structural diagram of the molecular compound to make sure the H is actually bonded to the N, O or F

21. O O O O H H H H H H H H                 • the hydrogen has such a low electroneg. in comparison to N, O and F so it has its electrons pulled so far away from it. This makes it able to be attracted not only to the  pole but also to the lone pairs

22. Other melting/ boiling point’s • Remember intermolecular forces only occur between molecular compounds and are weaker forces then intramolecular forces (bonds) • So when melting molecular compounds only the LD, DD, and HB need to be overcome • Metals and ionic compounds are attracted to one another by the bonds holding them together • Metallic structures and ionic compounds therefore have high bp/ mp due to having to overcome their intramolecular forces (bonding)

23. MP of Metals? • Metals are solid at room temp. because metal atoms have very strong forces between them • I.e. metallic bonding • So in order to melt them you need to add LOTS of energy (high temp) to overcome these strong forces

24. Metallic Bond Model metal cations “sea” of delocalized electrons

25. Ionic Compounds • Ionic compounds are also attracted to one another by strong forces (not quite a strong as metallic though) • I.e. ionic crystals • So in order to melt them you need to add quite a bit of energy (high temp) to overcome these forces

26. Ionic Crystals • ionic compounds have crystal structure oppositely charged ions • they form so that are asas possible close together • this is called a 3-D array of alternating positive and negative ions crystal lattice

27. Scale of Forces very high very low LD DD network covalent HB covalent ionic Intermolecular Forces (between) Intramolecular Forces (within) London Dispersion metallic ** wide range Dipole – Dipole ionic Hydrogen Bonding covalent network covalent eg) diamond, SiC, SiO2

28. Order of bp’s • Using the scale of forces you can order compounds based on their relative bp’s • Ex. From Highest to Lowest • Network covalent compound (ex. SiO2) • Ionic compound • Molecular compound with HB, DD, LD • Molecular compound with DD, LD • Molecular compound with LD (if 2 molecular compounds have LD only then bigger molecule or molecule with more electrons has higher bp)

29. Gases Unit Remember your formulas!!!!!

31. When to use what? • Use Boyle’s Law when temp. is constant • P1V1 = P2V2 • Use Charles’ Law when pressure is constant • V1/T1 = V2/T2 • Use Combined Gas Law when all three variables change • P1V1/T1 = P2V2/T2 • Use PV=nRT if given a mass or number of moles

32. Some others • If you are only given info about pressure and tempurature and its in a sealed container then V1 = V2 so using Combined Gas Law cancel out the volumes to get left with: • P1V1/T1 = P2V2/T2 • P1/T1 = P2/V2 • Law of Combining Volumes • You can use a balanced equation and multiply by coefficients wanted/given to get the volume of one gas if you know the volume of another

33. What volume of oxygen is used up if 100 mL of steam is formed in a composition reaction? What you are given What are  solving for O2(g) + 2H2(g) 2H2O(g)  x mL 100 mL 100 mLx 1 2 x mL = 50.0 mL

34. Convert 650 mmHg to kPa. Ratio of what you are trying to find Ratio of known values 101.325 kPa = x 760 mmHg 650 mmHg x = 86.6… kPa

35. Solutions, Acids and Bases

36. Remember your formula’s here too: • c=n/v • n=m/M • V1C1=VfCf • pH=-log[H30+] • pOH=-log[OH-] • [H30+]=10^-pH • [OH-]=10^-pOH • pH+pOH=14 Remember this is the dilution formula

37. Experiments • Remember in experiments there are always three variables: • Manipulated what you are changing • Responding the response to the change • Controlled what you keep the same

38. Ex: What effect does eating carrots have on eyesight? • Manipulated: amount of carrots eaten • Responding: how well you can see • Controlled: Same types of carrots, not eating any other food that could effect eyesight

39. Electrolytes? • Compounds that conduct electricity in water because they break apart into ions (ex. ionic compounds, acids) • Ex. NaCl  Na+(aq) + Cl-(aq) • Molecular compounds DO NOT break down into ions so they are non-electrolytes

40. Solubility • The ability to dissolve • If the solution is holding as many solutes as possible the solution is SATURATED and adding anymore solute will NOT be able to dissolve • A saturated solution usually has a small amount of UNDISSOLVED solute at the bottom. This is in constant EQUILIBRIUM with the solute that is dissolved in the solution (they switch places with each other all the time) Dissolved Equilibrium Undissolved

41. Standard Solution • Remember how to prepare a standard solution • Use formula’s n=m/M and c=n/v to get the mass you need for the certain volume and concentration • Steps: • Weigh out solute • Dissolve in ~half amount of water in a beaker • Pour into final volume volumetric flask • Fill flask, and invert to mix

42. Dilution • When you have a solution that has too high of a concentration you can add water to dilute it (water it down so its not as strong) • Ex. You have 100mL of a 5.0 mol/L solution. You add 500mL of water. What is the new concentration? Ci = 5.0 mol/L Vi = 0.1L Vf = 0.6L Cf = ? ViCi=VfCf Solve for Cf

43. Dissociation and dissociation equations • This is the same as ‘dissolving’ • When you have a compound and put it in water 4 situation to know: • It doesn’t dissolve • Ex. C25H52(s)  C25H52(s) • It does dissolve and its ionic • Ex. Ca(OH)2(s)  Ca2+(aq) + 2OH-(aq) • It does dissolve and its molecular • Ex. C12H22O11(s)  C12H22O11(aq) • It does dissolve and its an acid(this is a special case because it is a molecular compound but it acts as an ionic compound) • Ex. H2SO4(aq)  2H+(aq) + SO42-(aq)

44. Concentration of Ions • If asked to calculate the concentration of an ion in a solution 1st write the dissociation equation then treat it like a solution stoich. Question (no volumes are needed since they all have the same volume so you don’t need to calculate n first) • Ex. Calculate the ion concentrations when you have 0.500 mol/L H2SO4(aq) ? g w w H2SO4(aq)  2H+(aq) + SO42-(aq) c=0.500 mol/L c= ? c=? c of H+is = 0.500 mol/L x 2/1 = 1.00 mol/L c of SO42- is = 0.500 mol/L x 1/1 = 0.500 mol/L

45. Remember your properties of acids/ bases

46. Neutral Substances Acids Bases sour taste bitter taste electrolytes electrolytes electrolytes, non-electrolytes neutralize acids neutralize bases react with indicators do not react with indicators affect indicators the same way litmus - red litmus - blue bromothymol blue - yellow bromothymol blue - blue phenolphthalein - phenolphthalein - pink colourless react with to produce metals H2(g) pH greater than 7 pH of 7 pH less than 7 eg) HCl(aq), H2SO4(aq) eg) eg) Ba(OH)2(aq) NH3(aq) NaCl(aq), Pb(NO3)2(aq)

47. What is an acid? • TheArrhenius definition of an acid is that it has H+ at the beginning of the compound and is (aq) • Ex. HF(aq) • The modified Arrhenius definition of an acid is it reacts with water to form H3O+ ions • Ex. HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)

48. What is a base? • The Arrhenius definition of a base is that it has OH- ions at the end of an IONIC compound • Ex. NaOH(aq) • The modified Arrhenius definition of a base is that it reacts with water to form OH- ions. • Ex. Na2CO3(aq) + HOH(l)  NaOH(aq) + H2CO3(aq) OH- ions

49. Strong acids/ bases • Weak acids and bases don’t 100% break down to form H3O+ ions and OH- ions (strong one’s DO) • Strong acids are listed on the back of your periodic table • If they are NOT on that list they are a weak acid • Strong bases have OH- ions in them OR a metal with oxygen • Ex. NaOH and MgO • Every other base is a weak base • Ex. NH3 and Na2CO3

50. Monoprotic vs. Polyprotic • Monoprotic acids only have 1 H+ ion to give away (to water) • Ex. HCl, HF • Similarly monoprotic bases can only accept one H+ ion (from water) or has 1 OH- ion • Ex. NaOH, ions with only 1- charge (ex. F-) • Polyprotic acids have more than 1 H+ to give away • Ex. H2SO4, H3PO4 • Similarly polyprotic bases can accept more than 1 H+ ion • Ex. compound with ions that have more then 1- charge (ex. CO32-, PO43-)