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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. 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**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**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**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**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.**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**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**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**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**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**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**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**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**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**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**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**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**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**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**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**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**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**(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)**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**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**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**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**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**8. Nuclear Stability**Belt of Stability (Figure 1) Figure 1: Stable and unstable nuclides**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**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**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**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:**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**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-)**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**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**mass**charge B. Nuclear equations: Must obey Law of Conservation of ______ (top line) and Law of Conservation of _______ (bottom line). 1. Alpha emission (__________)**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-**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**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**np + e-**β - emission α - emission pn + e+ e- capture β + emission D. Nuclear Stability 1. Belt of Stability (Figure 1) Figure 1: Stable and unstable nuclides**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)**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**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

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