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Chemistry. First Semester Final Review By GowerHour, Inc. Unit 1: Algebra in Chemistry. Algebra in Chemistry: “What you do to one side of the equation, you must do to the _______.” 1. D = m/V is the given equation for the density of an object.

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    1. Chemistry First Semester Final Review By GowerHour, Inc.

    2. Unit 1: Algebra in Chemistry • Algebra in Chemistry: “What you do to one side of the equation, you must do to the _______.” 1. D = m/V is the given equation for the density of an object. If the mass and density of an object are given, what would be the equation to find the volume? • Solve E = mc2 for m? other V = m/D m = E/c2

    3. substitute 4. It is often required to _________one or more equations into another equation to form the working equation. • If the radius, mass, and density of a cylinder are known, the equation for the height of the cylinder can be formed by combining and the rearranging the density (D = m/V) and volume (V = r2h) formulae as follows: h = m/ D  r2

    4. Significant Figures • Any digit that is not zero is significant • 1.234 kg 4 significant figures • Zeros between nonzero digits are significant • 606 m 3 significant figures • Zeros to the left of the first nonzero digit are not significant • 0.08 L 1 significant figure • If a number is greater than 1, then all zeros to the right of the decimal point are significant • 2.0 mg 2 significant figures • If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant • 0.00420 g 3 significant figures

    5. How many significant figures are in each of the following measurements? 24 mL 2 significant figures 4 significant figures 3001 g 0.0320 m3 3 significant figures 6.4 x 104 molecules 2 significant figures 560 kg 2 significant figures

    6. 89.332 + 1.1 one significant figure after decimal point two significant figures after decimal point 90.432 round off to 90.4 round off to 0.79 3.70 -2.9133 0.7867 Significant Figures Addition or Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers.

    7. 3 sig figs round to 3 sig figs 2 sig figs round to 2 sig figs Significant Figures Multiplication or Division The number of significant figures in the result is set by the original number that has the smallest number of significant figures 4.51 x 3.6666 = 16.536366 = 16.5 6.8 ÷ 112.04 = 0.0606926 = 0.061

    8. 6.64 + 6.68 + 6.70 = 6.67333 = 6.67 = 7 3 Significant Figures Exact Numbers Numbers from definitions or numbers of objects are considered to have an infinite number of significant figures The average of three measured lengths; 6.64, 6.68 and 6.70? Because 3 is an exact number

    9. The number of atoms in 12 g of carbon: 602,200,000,000,000,000,000,000 The mass of a single carbon atom in grams: 0.0000000000000000000000199 Scientific Notation 6.022 x 1023 1.99 x 10-23 N x 10n N is a number between 1 and 10 n is a positive or negative integer

    10. move decimal left move decimal right Scientific Notation 568.762 0.00000772 n > 0 n < 0 568.762 = 5.68762 x 102 0.00000772 = 7.72 x 10-6 Addition or Subtraction • Write each quantity with the same exponent n • Combine N1 and N2 • The exponent, n, remains the same 4.31 x 104 + 3.9 x 103 = 4.31 x 104 + 0.39 x 104 = 4.70 x 104

    11. Scientific Notation Multiplication (4.0 x 10-5) x (7.0 x 103) = (4.0 x 7.0) x (10-5+3) = 28 x 10-2 = 2.8 x 10-1 • Multiply N1 and N2 • Add exponents n1and n2 Division 8.5 x 104÷ 5.0 x 109 = (8.5 ÷ 5.0) x 104-9 = 1.7 x 10-5 • Divide N1 and N2 • Subtract exponents n1and n2

    12. Converting within metric system using dimensional analysis 1. Convert to base unit by canceling units (Top unit cancels with bottom unit). 2. Place the multiplier with the base unit. 3. Place a 1 in front of the unit with prefix. 4. To enter multiplier into the calculator, use a 1 before the exponent key. Example: 10 -6 = 1 EE/EXP - 6 5. Metric dimensional analysis examples: a. Convert 3.6 nm to m. b. Convert 0.456 dag to pg. 0.456 dag 10 1 g 1 pg = 4.56 x 10 12 pg 1 dag 10 -12 g 3.6 nm 10 -9 m = 3.6 x 10 -9 m 1 nm

    13. 60 min m x x x 343 60 s 1 mi s 1 hour = 767 1 min 1609 m mi hour The speed of sound in air is about 343 m/s. What is this speed in miles per hour? meters to miles seconds to hours 1 mi = 1609 m 1 min = 60 s 1 hour = 60 min

    14. mass density = volume A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass? m m d = d = V V Unit 2 Density – SI derived unit for density is kg/m3 1 g/cm3 = 1 g/mL = 1000 kg/m3 = 21.5 g/cm3 x 4.49 cm3 = 96.5 g m = d x V

    15. I. Atomic Theory Unit 3: Nuclear Chemistry “_______ is composed of tiny _________ called _______.” Chart: Matter particles atoms Matter compound elements atoms nucleus electrons p+ n 3 quarks 3 quarks

    16. I. Historical Perspective: A. Early Ideas 1. Democritus (400 B.C.) a. ________ philosopher … b. Matter is composed of __________ _________ called atoms. c. “atomos” means ___________. d. His theory was forgotten because … 1) ______ and _________ disagreed. • He had no ______________ _________ • a) He used ________ not ____________. Greek indivisible particles indivisible Aristotle Plato experimental evidence theory experiments

    17. life gold 2. Alchemy a. 2 goals: 1) Elixir of ____. 2) Transmutation: Lead into ___________. b. Importance of alchemists is that they shifted from ________ to observation and _______________! 3. 1500’s – 1700’s: With the development of true _________ methods, scientists discovered many important things about matter. a. Such as: electricity, magnetism, chemical reactions; this information would help establish important scientific principles that would be used to develop the _______ Theory. experimentation thought scientific Atomic

    18. John theory behavior 4. Dalton’s Atomic Theory (1808) • ______ Dalton re-proposes the atomic _______ and supports his ideas with chemical ________. b. 3 Important Laws Dalton based his Atomic Theory: 1) Law of ____________ of Mass: Mass is neither ________ nor _________. 2) Law of Constant Proportions (________): Compound always contains the same elements in the same ___________ by mass. Example: H2O 3) Law of Multiple Proportions (___________): When two elements can form multiple compounds, the ratio of masses will remain constant for each compound. Example: CO : CO2 Conservation created destroyed Definite proportions John Dalton

    19. 1. Postulates: a. Elements are composed of small, indivisible particles called __________. b. Each element is made up of atoms that are __________ to each other. c. Chemical reactions are simple rearrangements of atoms in small, _____________ratios. 2. Problems with the Atomic Theory? • _________________________ • ______________________________________________________________________________ atoms identical whole number Atoms are divisible. Atoms of an element are not necessarily identical because of isotopes

    20. J.J. particles smaller Cathode Ray Tube • Thomson Model of the Atom Theory a. ____ Thomson discovered that atoms are made of ________, in other words they are made of _______ things. b. J.J. did experiments with ___________________ (CRT) and he found that: 1) Cathode rays are ___ particles that he called _________. 2) This showed that atoms are not __________! electrons indivisible

    21. Cathode Ray Tube

    22. charge to mass 3) He determined the _______________ ratio of the electron. 4) He knew atoms were _________ so he proposed a model of the atom called the ___________________. The plums were __. 5) Robert A. ________ determined the charge of the ___ with his oil ______ experiment. neutral Plum Pudding Model e- Millikan e- drop

    23. Measured mass of e- (1923 Nobel Prize in Physics) e-charge = -1.60 x 10-19 C Thomson’s charge/mass of e- = -1.76 x 108 C/g e- mass = 9.10 x 10-28 g

    24. radioactivity Rutherford 6. Rutherford Model of the Atom Theory a. The discovery of ____________ led to further advances in the atomic theory. b. New Zealand physicist Ernest B. __________ and his associates (Geiger & Marsden) used radioactive ___ particles to probe the _____. c. He discovered the ______________ with the ______ scattering experiment. d. Diagram of the experiment: (See Above) e. His calculations regarding the deflected particles indicated that atoms have a _____, __________, __________________. f. If the atom were a _____ in diameter, the nucleus would be the size of a ________…yet the nucleus contains virtually all of the atom’s ______. In other words most of the atom is made up of ______ _______.  atom alpha helium nucleus tiny + charged “massive” nucleus mile baseball mass empty space

    25. (1908 Nobel Prize in Chemistry) • particle velocity ~ 1.4 x 107 m/s (~5% speed of light) • atoms positive charge is concentrated in the nucleus • proton (p) has opposite (+) charge of electron (-) • mass of p is 1840 x mass of e- (1.67 x 10-24 g)

    26. orbit nucleus planets sun g. He proposed the still popular (yet wrong) planetary model: 1) electrons ______ the positive ________. 2) like the ________ around the _____. • Problems with a planetary atomic model; It could not explain… 1) Electron ________: classical physics theory says that a charged particle (like an electron) moving in a circular orbit would ______ energy and slow down, eventually collapsing out of its ______ and _______ into the ________. 2) Periodic _________ behavior. 3) Atomic _____ spectra. 7. The _____ model of the atom: proposed by Danish physicist ___________ would first explain (sort of) some of these problems. collapse lose orbit crash nucleus chemical line Bohr Neils Bohr Only the H spectrum

    27. 3 u u (Proton) d II. Atomic Structure: nucleus A. 3 particles 1. Proton (p) a. located in the ________ b. unit charge = __ c. mass = 1.67265 x 10-24 g (Proton) d. relative mass = __ (relative to the other particles) e. charge = + 1.6022 x 10-19 C (Coulomb) f. the mass is 1,836 x the mass of a _______. g. discovered by ___________ h. made of __ quarks: + 1 1 electron Rutherford

    28. Thomson energy -1 2. Electron (e) a. located outside the nucleus in ________ levels (shells). b. unit charge = ___ c. mass = 9.11 x 10-28 g (9.10953 x 10-28 g) d. relative mass = __ (tiny compared to n & p) e. charge = 1.6022 x 10-19 C (Coulomb) f. the mass is 1/1,836 the mass of a _______. g. discovered by __________. 0 proton

    29. Chadwick 3 u d (Neutron) d nucleus 0 3. Neutron (n) a. located in the _______. b. unit charge = __ “neutral” c. mass = 1.67495 x 10-24 g slightly more massive than a proton. d. relative mass = __ e. discovered by __________ in 1932 (Why so late? ___________________________) f. made of ___ quarks: 1 No charge – Harder to detect

    30. core nucleon Nuclide B. Nucleus 1. The central ____ of the atom. 2. Contains the neutrons and protons: _______ = a particle in the nucleus (n or p) 3. _______ = the nucleus of an atom. 4. Contains almost all the _________ of the atom. 5. Has a _____ volume compared to the _____ atomic volume. (Ping pong ball in the ___________). 6. Density of nucleus = _________ g/cm3! 7. Radionuclide: an unstable ________. Why some nuclei stable and others are unstable? (__________) 8. Nuclear stability is due to… a. Nuclear binding __________________ Holds protons and neutrons together. b. n/p ratio = stable atoms have a favorable n/p ratio. Examples: mass small total Astro dome 1013-1014 nucleus radioactive Forces (strong force) Larger atoms Smaller atoms n=p n>p

    31. 8. Nuclear Stability Belt of Stability (Figure 1) Figure 1: Stable and unstable nuclides

    32. protons III. Mass Relationships in atoms A. Atomic Number (Z) 1. Equal to the number of _______. a. Equals the number of ________ in a _____________. 1) Example: a) Oxygen b) Oxide ion (ions: charged particles) b. Each type of ________ has a specific number of protons…this determines the element’s ________. neutral atom electrons 8 p+ O 8 e- 8 p+ O2- 10 e- element identity

    33. Mass # Atomic # protons neutrons B. Mass Number (A) 1. Equals the number of _______ + the number of _________. a. It is the number of _________. 2. # of neutrons = ____________________________________ 3. Correct notation: • Nuclear Symbol To determine the number of neutrons: _____________ nucleons mass # − atomic # (n = A − Z) X A Charge X = generic element Z A - Z

    34. A X Mass Number Element Symbol Z Atomic Number 1 3 2 H (D) H (T) H 1 1 1 235 238 U U 92 92 Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei

    35. Substance Symbol Z A p+ e- n Potassium 40 26 30 Cl 37 D. Example: (a) Fill in the blanks. Use your periodic table. K 19 19 19 21 Iron Fe 56 26 26 Chlorine 17 17 20 17 (b)Write the nuclear symbols of the above elements:

    36. Ion name Ion symbol Gained/ lost e- A p+ e- n Aluminum ion lost 3 e- 27 Cu+ 35 Sulfide gained 2 e- 32 gained 1 e- 35 44 Copper ion Lost 1e- 64 29 28 lost 4 e- 50 68 Bromide Br- 79 36 Tin ion Sn 4+ 118 46 III. Atomic charged electrons E. Ion: A _________ atom. One that has gained or lost ___________. 1. cation: A ____________ charged atom (_______ electrons). 2. anion: A _____________ charged atom (_________ electrons). 3. Example: Ca2+ (calcium ion) F- (fluoride) 4. Example: Fill in the blanks. Use your periodic table. positively lost negatively gained Lost 2e- Gained 1e- 20p+, 18e- 9 p+, 10e- Al3+ 13 10 14 S 2- 16 18 16

    37. atomic mass protons neutrons F. Isotopes: Atoms with the same _______ number but different _____ numbers. 1. Same number of ___________ 2. Different number of _____________. 3. Isotopes are named by their _______. • Carbon-14 b. Uranium-238 4. Example: The three isotopes of Hydrogen 5. Example: The three oxygen isotopes 6. Ions: ________ atoms (or groups of atoms) that have lost or gained electrons. Examples: mass # charged (gained 2 e-) (lost 3 e-)

    38. 14 11 C C 6 6 How many protons, neutrons, and electrons are in How many protons, neutrons, and electrons are in ? ? Do You Understand Isotopes? 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons

    39. p+ IV. Nuclear Chemistry: Chemistry of the nucleus (___ + ___). n spontaneous particles Electromagnetic radiation A. Radioactivity: The ____________ emission of _________ or EMR (_______________________) from the nucleus. 1. Henri Becquerel: Discovered radioactivity (1896) using ___________________ and uranium ore. 2. Types of Radioactivity: a. Alpha (_____): Nucleus of a helium atom. (_________) b. Beta (_____): High speed electron emission from the ________. (_________) c. Gamma ray (______): Photon of high energy light. (____) 3. Penetrating power of radiation: a. Alpha can go through ______. b. Beta particles can go through _________. c. Gamma rays can go through _______________________. ________>________>________ photographic plates nucleus paper 3 mm Al 3 cm Pb

    40. mass charge B. Nuclear equations: Must obey Law of Conservation of ______ (top line) and Law of Conservation of _______ (bottom line). 1. Alpha emission (__________)

    41. nucleus neutrons 2. Beta emission (________or________) a. Electron is formed in the ________. b. Nucleus has too many _________. c. Neutron spontaneously becomes a _________, which causes a high energy electron to be ejected from the nucleus. proton 6 p+ 7 p+ 8 n 7 n Net: n  p + e-

    42. excited 3. Gamma emission (_________) a. Nucleus is in an _______ energy state (excess energy from another decay). b. As nucleus loses energy, a ____________ is emitted. c. Asterisk (*) is used to symbolize ___________________. gamma ray excited energy level

    43. C. Decay Series: daughter A radioactive decay often results in a ___________ nucleus that is also radioactive. A radioactive ________ ______ refers to successive decays which starts with one parent isotope and proceeds through a number of daughter isotopes. The series ends when a stable, ______________ isotope is produced. decay series non-radioactive

    44. np + e- β - emission α - emission pn + e+ e- capture β + emission D. Nuclear Stability 1. Belt of Stability (Figure 1) Figure 1: Stable and unstable nuclides

    45. Figure2: A plot of nuclear binding energy per nucleon versus mass number. Nuclei with large binding energies per nuclei are the most stable. 2. Binding Energy per Nucleon versus Atomic Mass (Figure 2)

    46. 3. Why are some nuclei stable and others are not? Rule 1: The greater the binding energy per nucleon (energy holding the nucleus together), the greater the stability. (See Figure 2). Which isotope is the most stable, according to Fig. 2? ______ Rule 2: Nuclei of low atomic numbers with a 1:1 ratio of neutrons to protons are very stable. (See Figure 1) Example: Carbon-12 Helium-4 ** Radioactive isotopes decay until they reach the “Belt of Stability.” Rule 3: The most stable nuclei tend to contain an ______ number of both protons and neutrons. Example: Iron-56 Oxygen-16 2 p+, 2n 6 p+, 6n even 26 p+, 30n 8 p+, 8n

    47. neutrons protons E. How is Binding Energy determined? 1. Nuclear binding energy: The energy required to break up a nucleus into its component _______ and ________ (nucleons). 2. Binding energy comes from the mass defect of the nucleus. 3. Mass Defect: The total mass of the stable nucleus _________ the sum of the masses of the nucleons. The “missing” mass has converted into energy! (_______) Masses of subatomic particles e- 0.00054858 amu p+ 1.007276 amu n 1.008665 amu (p+ + e-) 1.007825 amu less than E = mc2 Note: When calculating masses, include the mass of the electrons because the mass of the whole atom includes electrons. NOTE: The whole atom is NOT the sum of its parts. MISSING MASS Mass defect = Atomic mass of the isotope -  mass of subatomic particles. Relationship between mass and energy: E = mc2