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Final Exam Study Notes

Final Exam Study Notes. CP Chemistry Period 5. Physical and Chemical Properties. Physical Property- a trait or characteristic that you can observe without changing the identity of the substance Chemical Property - a trait you can observe by changing the identity of the substance.

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Final Exam Study Notes

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  1. Final Exam Study Notes CP Chemistry Period 5

  2. Physical and Chemical Properties Physical Property- a trait or characteristic that you can observe without changing the identity of the substance Chemical Property -a trait you can observe by changing the identity of the substance

  3. Physical and Chemical Changes Physical Change- a change that affects only physical properties and does not alter the identity of the substance Chemical Change- a change that alters the identity of our substance

  4. Elements, Compounds, and Mixtures • Element- a substance that cannot be separated by chemical or physical means • Compound- a substance made up of two or more elements only separated by chemical means • Atom- smallest unit of an element • Molecule- smallest unit of a compound

  5. Elements, Compounds and Mixtures cont. • Mixture- a combination of substances thhat are not chemically combined. These can be separated physically • Homogeneous: looks same throughout • Heterogeneous: composed of different parts

  6. Accuracy and Precision True Data is accurate. Repeatable data is precise. Accurate & Precise

  7. Counting Protons/Neutrons/Electrons • The mass number of the element eqauls the number of protons in an element • The number of protons is also the number of electrons unless its an ion • To find the neutrons you subtract the atomic number from the mass number

  8. Isotopes • Are atoms that have lost or gained neutrons, same element but different number of neutrons

  9. Radiation particles/ Nuclear equations • Alpha- 42H- stopped by clothing or skin • Beta- 0 -1 e- stopped by a sheet of lead • Gamma- stopped by several inches of lead, most dangerous • Nuclear reactions happened when there is an unstable particle and eventually gives off a particle of radiation

  10. Ions • Atoms that have lost or gained electrons • atoms turn into ions when electrons move • Ions have a charge • There are negative electrons and positive electrons • How Ions are formed: • Positive ions have lost electrons • Negative ions have gained electrons • Positive ions are called cations • Negative ions are called anions • When atoms are most stable they have an octet • Octet- 8 electrons in the outer most energy level

  11. Covalent Bonds • Share electrons between two atoms • Properties • Low melting point • Molecule structure • Gases, liquids, soft solids • Poor conductors of heat • Poor conductors of electricity • Typically 2 non-metal atoms

  12. Ionic Bonds • They trade electrons between the two atoms • Ions must form from the atom • Properties • High melting point • Crystal lattice structure • Hard solids • Brittle • Good conductors of heat • Good conductors of electricity • Typically 1 metal and 1 non metal

  13. Wavelength • Waves of Light: electromagnetic radiation (light) moves as a wave Crest Wave Trough

  14. Wavelength/Frequency • λ= Wavelength: Distance from crest to crest on a wave • v= Frequency: How often a wave passes by in a second (s-1) • Wavelength and Frequency are inversely related • Wavelength increases, Frequency decreases • Wavelength decreases, Frequency increases

  15. Calculations • E=h • E= Energy (Joules) • h= Plank’s Constant (6.626 x10-34) • = Frequency • Example: A yellow light has a wavelength of 600nm • a) What is the frequency of the light? • b) What is the Energy of the light? • Answers on next slide

  16. Answer – Frequency • a) = 600nmn v= x c= 3.0 x 108 m/s = c/v 600nm x 1n/10 x 9nm= 6.0 x 10-7m 6.0 x 10-7m= 3.0 x 108 m/s /v V= 3.0 x 108 m/s / 6.0 x 10-7m = 5.0 x 1014 s-1

  17. Answer- Energy • E=hv • E= ? • h= 6.626 x10-34 J(s) • v= 5.0 x 1014 s-1 • E= (6.626 x10-34 ) 5.0 x 1014 • 3.31 x 10-19 J

  18. Ionization Energy • The amount of energy needed to remove one electron from an atom • Size of atom determines how easily electrons are removed • Big atoms lose e- with minimal effort • Little atoms lose e- with a huge amount of energy needed - Increases up and to the right on the Periodic Table • Noble gases all have elect negativities equal to zero • If an atom needs a lot of energy to remove an electron its because it really wants the e-

  19. Periodic Trends • Properties of elements can be predicted using the location on the periodic table • Electron Configuration • Family and Periods • Densities • Reactivity • Atomic radius • Ionization Energy • Electronegativity

  20. Atomic Radius • The distance from the nucleus to the outer-most elections (in the highest energy orbital filled) • As electrons fill into higher energy orbitals, the radius of the atoms gets bigger! Rb 5s1 K 4s1 Rb Na 3s1 K Na

  21. Atomic Radius • Within an energy level adding more protons makes the radius of atoms smaller because the protons can hold the electrons in closer P 3P3 S 3 P4 P Cl 3P5 S Cl

  22. Question • Put in order smallest to largest: • Rb, P, Na • Answer on next slide

  23. Answer • Na, Rb, P

  24. Periodic Table Families and Periods • Groups (families) = The columns on the Periodic Table • Periods = Rows on the Periodic Table • Elements arranged by atomic number • Column 1: #1 • Silvery White • Column 2: • React with water to form a base • Column 3: 5,8,9,10,11,12 • All metals • Form colorful solutions • Hard Brittle • Metallic • Versatile in bonding ability • Charges: +2, +3

  25. Periodic Table Families and Periods • Column 17: Halogens • Non-metallic • Charge -1 • Very reactive with positive ions • Not solid at room temperature • Colored Gas • Column 18: Noble Gases • Non-Metallic • Non Reactive • Charge: 0 • Column 14 • Charge +4 -4 • Non-metallic • Can bond with positive or negative ions • Solid at room temperature • Relatively low reactivity • Column 16: Oxygen Family • Charge: -2 • Non-metallic • Bonds with itself • Shares electrons with other elements

  26. VSEPR Molecule Geometries • VSEPR is a model of molecular structures based on the idea that ideal structures minimize electron pair repulsions • Used to draw and evaluate Lewis Structures • Bare electrons are the most repulsive!

  27. Molecular Geometry Models • Looking at the molecular geometry of a single atom, not of an entire molecule • 3D Figures to represent Lewis Structures • Constituent groups are the things bonded to the atom under scrutiny • Dashed lines represent a bond behind the plane of the paper; wedged lines represent a bond coming toward you

  28. Planar Geometry • Linear • 1-2 Constituents • 0 Lone Pairs • Bond Angle: 180 • Trigonal Planar • 3 Constituents • 0 Lone Pair • Bond Angle: 120 • Bent • 2 Constituents • 1 Lone Pair • Bond Angle: <120

  29. Tetrahedral & Derivatives • Tetrahedral • 4 Constituents • 0 Lone Pair • Bond Angle: 109.5 • Trigonal Pyramidal • 3 Constituents • 1 Lone Pair • Bond Angle: 107.3 • Bent • 2 Constituents • 2 Lone Pairs • Bond Angle: 104.5

  30. Lewis Dot Diagrams • Lewis dot structure is a drawing of how the atoms are bonded together covalently using valence electrons. • You need to know • Shared pair= 2 electrons shared by 2 atoms (bond) • Lone pair= 2 electrons not shared by atoms (unshared pair)

  31. How to draw Lewis Dot Structure 1. Count the total valence electrons for the molecule Ex: SCl2=20 valence electrons 2. Select a central atom. look for= *the only one of its kind. *less electronegative Ex: SCl2= S is central atom because it’s alone

  32. How to draw Lewis Dot Structure 3. Set up the elements as symmetrical as possible. Ex: SCl2= Cl S Cl 4. Draw in shared pairs by drawing a line. Ex: SCl2= Cl-S-Cl

  33. How to draw Lewis Dot Structure 5. Account for electrons used from total you started with.(Shared pairs=2 electrons) Ex: 20 valence electrons -4 shared electrons ___ 16 unshared electrons

  34. How to draw Lewis Dot Structure 6. Fill in unshared pairs around the outside of elements Ex:

  35. How to draw Lewis Dot Structure 7. When there isn’t enough electrons for every element to have an octet, we share more pairs. Ex:

  36. How to find polarity in molecules • When finding the polarity in molecules you need to find out if the bonds are polar or non-polar first. • Polar bond- when 2 atoms share electrons unequally • Non-Polar bond- share electrons completly even.

  37. How to find polarity in molecules • Polar Molecules- Molecules where one side of the molecule has more electrons than the other. • 1. if there is a lone pair on the center atom, it is polar • 2. If bonds are unequal polarity, than the molecule is polar

  38. How to find polarity in molecules • Non-Polar Molecules • If there is no lone pairs on the center atom • If the bonds are equal polarity

  39. 5 Types of Chemical Reactions • Synthesis: • Ex: Cu+3 + O2 Cu2O3 • Decomposition • Ex: MgCl2 Mg+2 + Cl • Single Replacement • Ex: MgCl2 + Cu+2  Mg + CuCl2 • Double Replacement • Ex: 3MgCl2 + Cu2O3  3MgO + 2CuCl3 • Combustion • Hydrocarbon + oxygen  CO2 +H2O • CH3OH + O2  CO2 + H2O

  40. Predicting Products (using the 5 type of chemical reaction) • Synthesis all you have to do is combine the two reactants to get your products • Decomposition you break up your reactants and get two products • Single replacement take either the positive or negative ion (by itself) and replace it with the positive or negative ion from a formula in the equation • Double replacement you take the positive ion from one formula and put the negative ion from the other formula to create a new formula, do this again with the left over positive and negative ions. (positive come first) • Combustion always ends up with carbon dioxide and water

  41. Law of Conservation of Matter • Matter can neither be created nor destroyed • In chemical equation it’s crucial to make sure its balanced because, if its not balances it goes against the law of conservation of matter because it creates (or destroys) matter

  42. Balancing Equations • According to the law of conservation of matter you have to balance all of your equations so that you don’t create or destroy matter. • NaOH + Cl2 NaCl + OH • This is not balance because you have two chlorines in the reactants and only one on the product side • So, all you have to do is add a coefficient in order to balance it: • 2NaOH + Cl2  2NaCl + 2OH

  43. Net Ionic Equations • Aqueous substance dissociate • A complete ionic equation shows all the ions and molecules in a reaction • Zn (s) + CuSO4 (aq)  ZnSO4 (aq) • Complete ionic equation: • Zn (s) + Cu+2 (aq) + SO4-2 (aq)  Zn+2 (aq) + SO4-2 (aq) +Cu (s) • NET IONIC: • Anything that’s AQUEOUS that’s the same on both sides you can cancel out • So you can get rid of SO4-2 (reactant side) AND SO4-2 (product side) • Final net ionic equation: • Zn (s) + Cu (aq)  Zn (aq) + Cu (s)

  44. Problems • 1. What type of reaction is this? • Mg + KCl MgCl + K • 2. Balance the following equations: • ___ Al + ___ O2  _____ Al2O3 • ___CuS + ____ O2  ____CuO + _____ SO2 • ____ Ca3P2 + ____ H2O  ____ Ca(OH)2 + ___ PH3 • 3. Write the net ionic equation for: • AgNO3 (aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)

  45. ANSWERS • 1. single replacement • 2. 4 Al + 3O2  2 Al2O3 • 3. 2 CuS+ 3 O2  2 CuO+ 2 SO2 • 4. (1) Ca3P2+ 6 H2O  3 Ca(OH)2 + 2 PH3 • 5. AgNO3 (aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • Complete ionic = • Ag+1 (aq) + NO3-1 (aq) + Na+1(aq) + Cl-1(aq)  Ag+1 (aq) + Cl-1 (s) + Na+1 (aq) + NO3-1 (aq) • Net Ionic = • Ag+1 (aq) +Cl-1 (aq)  AgCl (s)

  46. The Factor Label Method • A ratio used to convert the unit you have into the desired unit • Example: If you are given one day, how do you convert it into the amount of seconds in a day? Answer: 1 day* 24 hours*60 minutes*60 sec 1 day 1 hour 1 min Cross cancel the units!!

  47. The Factor Label Method Cont. • The factor label method is useful in converting metric prefixes • The Metric Prefixes are: • TGMKHDBDCMMNP • The Great Mister King Henry Died By Drinking Chocolate Milk Monday Night Partying • Tetra Giga Mega Kilo HectoDeca BASE DeciCentiMili Micro Nano Pico You can use Metric conversions to change from prefix to prefix!

  48. Converting Moles, Liters, Grams & Particles • 1 Mole = 6.02 X 1023 “things” • 6.02 X 1023 = Avogadro’s Number • Moles to Particles/Liters/Grams: mole of element X 6.02*1023 = # with desired unit 1 Mole

  49. Molar Mass • To find the molar mass you must refer to the periodic table • Look up each atomic mass of the element and add them all together to find the molar mass • If there is a subscript then you must multiply the atomic mass by the subscript

  50. Empirical Formula • Empirical formula= the lowest terms ratio of elements in a formula (Not the true formula) • To calculate the empirical formula you must find • Percent to Mass • Mass to Mole • Divide by small • Multiply until whole

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