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Unit 3: “ Atomic Structure ” Chemistry: Mr. Blake/Mr. Gower.

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2 Unit 3: “ Atomic Structure ” Chemistry: Mr. Blake/Mr. Gower

3 I. Atomic Structure NOTE: The Greek philosopher __________ (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “ atomos ” ) –He believed that atoms were ________ and _____________ –His ideas did agree with later scientific theory, but did not explain chemical behavior, and was ________________ _______________– but just philosophy Democritus indivisible indestructible not based on the scientific method A. Section 4.1 Defining the Atom

4 1. Dalton ’ s Atomic Theory (experiment based!) c. Atoms of different elements combine in simple whole-number ratios to form chemical _________ d. In chemical reactions, atoms are combined, separated, or rearranged – but ____ changed into atoms of another element. a. All elements are composed of tiny indivisible particles called _____. b. Atoms of the same element are _______. Atoms of any one element are _______ from those of any other element. John Dalton (1766 – 1844) atoms identical different compounds never

5 Problems with Dalton’s Atomic Theory? 1. matter is composed of indivisible particles Atoms Can Be Divided, but only in a nuclear reaction 2. all atoms of a particular element are identical Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)! Different elements have different atoms. YES! 3. atoms combine in certain whole-number ratios YES! Called the Law of Definite Proportions 4. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements. Yes, except for nuclear reactions that can change atoms of one element to a different element

6 2. Sizing up the Atom a. Elements are able to be subdivided into smaller and smaller particles – these are the _____, and they still have __________ of that element b. If you could line up 100,000,000 copper atoms in a single file, they would be approximately ________ c. Despite their ________, individual atoms ___ observable with instruments such as scanning tunneling (electron) microscopes atoms properties 1 cm long small size are

7 B. Section 4.2 Structure of the Nuclear Atom NOTE: One change to Dalton ’ s atomic theory is that _______________ into subatomic particles: NOTE: ________________________ are examples of these fundamental particles NOTE: There are many other types of particles, but we will study these three atoms are divisible Electrons, protons, and neutrons

8 1. Discovery of the Electron a. In 1897, J.J. Thomson used a _______ _______ to deduce the presence of a negatively charged particle: the _______ cathode ray tube electron

9 Modern Cathode Ray Tubes  Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. TelevisionComputer Monitor

10 2. Mass of the Electron a – _____________ determines the _____ of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge The oil drop apparatus Mass of the electron is 9.11 x g Robert Millikan mass

11 3. Conclusions from the Study of the Electron: a. Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. b. Atoms are neutral, so there must be ______________ in the atom to balance the negative charge of the electrons c. _______________________ that atoms must contain other particles that account for most of the mass positive particles Electron have so little mass

12 d. ____________ in 1886 observed what is now called the “ ______ ” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron) e – _____________ confirmed the existence of the “ ______ ” – a particle with ________, but a mass nearly ____ to a proton Eugen Goldstein proton James Chadwick neutron no charge equal

13 4. Subatomic Particles ParticleChargeMass (g)Location Electron (e - ) 9.11 x Electron cloud Proton (p + ) x Nucleus Neutron (n o ) x Nucleus

14 Nucleus Electron cloud

15 5. Thomson ’ s Atomic Model a. Thomson believed that the ________ were like plums embedded in a positively charged “ pudding, ” thus it was called the “ __________ ” model. J. J. Thomson electrons plum pudding

16 6. Ernest Rutherford ’ s Gold Foil Experiment a. Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil b. Particles that hit on the detecting screen (film) are recorded

17 7. Rutherford’s problem: a. In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target? Target #1 Target #2

18 b. The Answers: Target #1 Target #2

19 8. Rutherford ’ s Findings 1) The nucleus is _____ 2) The nucleus is _____ 3) The nucleus is _______ charged a.Most of the particles passed right through b.A few particles were deflected c. VERY FEW were greatly deflected “ Like howitzer shells bouncing off of tissue paper! ” d. Conclusions: small dense positively

20 9. The Rutherford Atomic Model a. Based on his experimental evidence: 1) The atom is mostly empty space 2) All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “ ______ ” 3) The nucleus is composed of ______ and ________ (they make the nucleus!) 4) The electrons distributed around the nucleus, and occupy most of the ______ 5) His model was called a “ ___________ ” nucleus protons neutrons volume nuclear model

21 1. Atomic Number a. Atoms are composed of _______ protons, neutrons, and electrons b. How then are atoms of one element different from another element? c. Elements are different because they contain different numbers of ________ d. The “ ____________ ” of an element is the _______________ in the nucleus e. ___________________________ identical PROTONS atomic number number of protons # protons in an atom = # electrons C. Section 4.3: Distinguishing Among Atoms

22 2. Definition: Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element# of protonsAtomic # (Z) Carbon Phosphorus Gold

23 3. Mass Number Definition: Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p + + n 0 Nuclidep+p+ n0n0 e-e- Mass # Oxygen Arsenic Phosphorus

24 4. Nuclear/Complete Symbols a. Contain the symbol of the element, the mass number and the atomic number. X Mass number Atomic number Subscript → Superscript →

25 a. Find each of these: 1) number of protons 2) number of neutrons 3) number of electrons 4) Atomic number 5) Mass Number Br Symbols

26 b. If an element has an atomic number of 34 and a mass number of 78, what is the: 1)number of protons 2)number of neutrons 3)number of electrons 4)complete symbol Se 78 34

27 d. If an element has 78 electrons and 117 neutrons what is the 1)Atomic number 2)Mass number 3) number of protons 4) complete symbol Pt

28 a. Dalton was wrong about all elements of the same type being identical b. Atoms of the same element can have different numbers of _______. c. Thus, different mass numbers. d. These are called _______. neutrons isotopes 6. Isotopes

29 Isotopes e. _____________( ) proposed the idea of isotopes in 1912 f. _______ are atoms of the ____ ______ having different masses, due to varying numbers of neutrons. g. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials. Frederick Soddy Isotopes same element

30 a. We can also put the mass number after the name of the element: b. Examples: carbon-12 carbon-14 uranium Naming Isotopes

31 c. _______ are atoms of the ___________ having ________ masses, due to varying numbers of neutrons. IsotopeProtonsElectronsNeutronsNucleus Hydrogen–1 (protium) 110 Hydrogen-2 (deuterium) 111 Hydrogen-3 (tritium) 112 Isotopes same element different

32 8. Isotopes a. Elements occur in nature as _______ of _______. b. Isotopes are atoms of the same element that differ in the _______ _______. mixtures isotopes number of neutrons

33 a. ____ are atoms or groups of atoms with a positive or negative charge. b. _________ an electron from an atom gives a _____ with a ____________ c. ______ an electron to an atom gives an _____ with a ____________. d. To tell the difference between an atom and an ion, look to see if there is a charge in the _________! Examples: Na + Ca +2 I - O -2 Na Ca I O Ions Taking away cation positive charge Adding anion negative charge superscript IONS 9. IONS

34 e. A cation forms when an atom loses one or more electrons. f. An anion forms when an atom gains one or more electrons Mg --> Mg e- F + e- --> F -

35 metals (Mg) lose electrons ---> cations nonmetals (F) gain electrons ---> nonmetals (F) gain electrons ---> anions NOTE: In General……

36 Learning Check – Counting State the number of protons, neutrons, and electrons in each of these ions. 39 K + 16 O -241 Ca #p + ___________________ #n o ___________________ #e - ___________________

37 One Last Learning Check Write the nuclear symbol form for the following atoms or ions: A. 8 p +, 8 n, 8 e - ___________ B.17p +, 20n, 17e - ___________ C. 47p +, 60 n, 46 e - ___________

38 Charges on Common Ions By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.

39 Example: A student has a test percentage of 78%; a lab percentage of 92%; and has completed homework at 100%. Her weighted average grade is computed as (78% X 0.6) + (92% X 0.20) + (100% X 0.20) = 84.5

40 a. How heavy is an atom of oxygen?  It depends, because there are different _____ of oxygen atoms. b. We are more concerned with the _________________. c. This is based on the abundance (percentage) of each variety of that element in nature. d. We don’t use grams for this mass because the numbers would be too small. kinds average atomic mass 10. Atomic Mass

41 a. Instead of grams, the unit we use is the ______________ (amu) b. It is defined as one-twelfth the mass of a carbon-12 atom. c. Carbon-12 chosen because of its _____ ______. d. Each isotope has its own atomic mass, thus we determine the average from percent abundance. Atomic Mass Unit isotope purity 11. Measuring Atomic Mass

42 a. Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. b. If not told otherwise, the mass of the isotope is expressed in _____________ (amu) atomic mass units 12. To calculate the average:

43 IsotopeSymbolComposition of the nucleus % in nature Carbon C6 protons 6 neutrons 98.89% Carbon C6 protons 7 neutrons 1.11% Carbon C6 protons 8 neutrons <0.01% Atomic mass is the average of all the naturally occurring isotopes of that element. Carbon = Atomic Masses

44 - Page 117 Question Solution Answer Knowns and Unknown

45 14. Calculate the atomic mass of carbon. a. Isotopes% abundance Atomic mass Carbon % amu Carbon % amu. b. Lithium has two isotopes. If lithium-6 has a mass of and 7.42 % occurrence, what is the % abundance and mass of lithium -7? Atomic mass= (%) (mass) +(%) (mass) +….. Atomic mass= (0.9889) (12.000) +(0.0111) (13.003) = = amu 6.941= (0.0742)(6.015) +(0.9258)(x) = amu

46 a. A “ periodic table ” is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties 1) The periodic table allows you to easily compare the properties of one element to another 15. The Periodic Table: A Preview

47 b. Each horizontal row (there are 7 of them) is called a _____ c. Each vertical column is called a ____________ 1) Elements in a _____ have similar chemical and physical properties 2) Identified with a number and either an “ A ” or “ B ” period group or family group

48 A few elements, such as gold and copper, have been known for thousands of years - since ancient times Yet, only about __ had been identified by the year As more were discovered, chemists realized they needed a way to ________ the elements. 13 organize II. The Periodic Table A. Section 6.1: Organizing the Elements NOTES:

49 Chemists used the _________ of elements to sort them into groups. In 1829 J. W. Dobereiner arranged elements into _____ – groups of three elements with similar properties One element in each triad had properties intermediate of the other two elements properties triads

50 a. By the mid-1800s, about 70 elements were known to exist b. Dmitri _________ – a Russian chemist and teacher c. Arranged elements in order of _________________ d. Thus, the first “ Periodic Table ” Mendeleev increasing atomic mass 1. Mendeleev ’ s Periodic Table

51 a. ___________ for yet undiscovered elements b. When they were discovered, he had made good predictions c. But, there were problems: –Such as Co and Ni; Ar and K; Te and I He left blanks 2. Mendeleev

52 a. In 1913, Henry ______ – British physicist, arranged elements according to increasing ____________ b. The arrangement used today c. The symbol, atomic number & mass are basic items included-textbook page 162 and 163 Moseley atomic number 3. A Better Arrangement

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54 a. When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. b. Horizontal rows = ______ 1) There are __ periods c. Vertical column = _____ (or family) 1) Similar physical & chemical prop. 2) Identified by number & letter (IA, IIA) periods 7 group 4. The Periodic Law

55 5. Areas of the Periodic Table Three classes of elements are: 1) _____, 2) ________, and 3) _________ 1) Metals: _______ conductors, have luster, ductile, malleable 2) Nonmetals: generally brittle and non-lustrous, poor conductors of ____ and electricity metalsnonmetals metalloids electrical heat

56 Areas of the Periodic Table Some nonmetals are _____ (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br) Notice the heavy, stair-step line? 3) _________: border the line-2 sides –Properties are __________ between metals and nonmetals gases Metalloids intermediate

57 Squares in the Periodic Table The periodic table displays the ______ and _____ of the elements, along with information about the structure of their atoms: Atomic ______ and atomic _____ Black symbol = solid; red = gas; ____ _____ (from the Periodic Table on our classroom wall) symbols names number mass blue = liquid

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59 Groups of Elements - Family Names Group IA (1) – __________ –Forms a “ base ” (or alkali) when _______ with water (not just dissolved!) Group 2A (2)– ________________ –Also form bases with water; do not dissolve well, hence “ earth metals ” Group 7A (17) – _______ –Means “ salt-forming ” Group 8A (18) – _________ –Nonreactive because of their electron configuration alkali metals alkaline earth metals halogens noble gases reacting

60 ELEMENTS THAT EXIST AS DIATOMIC MOLECULES Remember: HOFBrINCl These elements only exist as PAIRS. Note that when they combine to make compounds, they are no longer elements so they are no longer in pairs!

61 Nuclear Chemistry Chemistry – Unit 4 Chapter 25

62 Mass Defect Difference between the mass of an atom and the mass of its individual particles amu amu

63 Nuclear Binding Energy Energy released when a nucleus is formed from nucleons. High binding energy = stable nucleus. E = mc 2 E:energy (J) m:mass defect (kg) c:speed of light (3.00×10 8 m/s)

64 Nuclear Binding Energy Unstable nuclides are radioactive and undergo radioactive decay.

65 Types of Radiation Alpha particle (  ) –helium nucleus paper 2+ Beta particle (  -) –electron 1- lead Positron (  +) –positron 1+ Gamma (  ) –high-energy photon 0 concrete

66 Nuclear Decay Alpha Emission parent nuclide daughter nuclide alpha particle Numbers must balance!!

67 Nuclear Decay Beta Emission electron Positron Emission positron

68 Nuclear Decay Electron Capture electron Gamma Emission –Usually follows other types of decay. Transmutation –One element becomes another.

69 IQ# 1 1.Balance the following equations:

70 Nuclear Decay Why nuclides decay… –need stable ratio of neutrons to protons

71 Belt of Stability and Radioactive Decay

72 Half-life Half-life (t ½ ) –Time required for half the atoms of a radioactive nuclide to decay. –Shorter half-life = less stable.

73 Half-life m f :final mass m i :initial mass n:# of half-lives

74 Half-life Fluorine-21 has a half-life of 5.0 seconds. If you start with 25 g of fluorine-21, how many grams would remain after 60.0 s? GIVEN: t ½ = 5.0 s m i = 25 g m f = ? total time = 60.0 s n = 60.0s ÷ 5.0s =12 WORK : m f = m i (½) n m f = (25 g)(0.5) 12 m f = g

75 Example: How much of a 500. g sample of Uranium-235 would be left after five half-lives? (n = # of half-lives) M i = 500 g n = 5 M f = ? m f = m i (½) n m f = (500 g)(0.5) 5 m f = 15.6 g

76 Example: A mg sample of Radon- 222 decays to mg after 24 hours. Determine the half-life. 16 → 8 → 4 → 2 → 1 → 0.5 → = 6 half lives

77 Example: The half-life of molybdenum-99 is 67 hours. How much of a mg sample is left after 335 hours? M i = mg Half-life = 67 h Rxn time = 335 h M f = ? n = 335 / 67 = 5 m f = m i (½) n m f = (1.000 mg)(0.5) 5 m f = mg

78 Learning Check! The half life of I-123 is 13 hr. How much of a 64 mg sample of I-123 is left after 39 hours? M i = 64 mg n = 3 M f = ? m f = m i (½) n m f = (64 mg)(0.5) 3 m f = 8.0 mg

79 Half-life Lab Procedure: 1. Each lab group will acquire a sample of 50 pennies in a cup. 2. Count pennies to make sure you have 50 pennies. 3.Enter “50” in Shake # 0 row for Trial 1, 2, & 3 and “150” for  (Sum of) of trials Shake the cup of pennies. Pour the pennies on to the lab bench. 5. Remove all pennies that land on “heads”. They have decayed. 6. Count only the remaining pennies (the pennies that landed on “tails”). Record data. 7. Place only the remaining pennies (“tails”) into the cup and shake again. Repeat steps 4-7 until all pennies have decayed. 8. Repeat the process two more times and record data under Trial 2 & 3.

80 Data:Collect data for three trials in the table. Data Analysis: Prepare a graph to represent the decay of your sample (  of trials (y-axis) vs. Shake # (x-axis)) Prepare a graph in your lab book: Graph the # of undecayed atoms (  of trials) (y-axis) versus the Shake # (x-axis). Label the x and y axes, including units (if applicable). Make graph large (at least 2/3 pg.). Draw a best fit curve that represents your data. Use a Ruler! Plot the Shake # for the  of trials using the best fit curve. Determine the “half life” of your sample in terms of # of shakes using your graph.

81 Graphing the Results Important !! Title every graph and label each axis (include units) Graphs should be at least 2/3 page Use a ruler Circle all data points Use a best-fit line (no “connect the dots”!) Find the average half-life (in # of trials) of your sample by interpolating your curve at exactly 75, 37.5, and pennies undecayed)

82 Half-Life Lab Σ of Trials Shake # · · · · Use Ruler for axis Label Axis Best Fit Curve Convenient #’s Circle Data Points At least 2/3 of pg Title = ~ 1.1 shake Half-life = 1.1 shake 123

83 F ission splitting a nucleus into two or more smaller nuclei 1 g of 235 U = 3 tons of coal

84 F ission chain reaction - self-propagating reaction critical mass - mass required to sustain a chain reaction

85 Fusion combining of two nuclei to form one nucleus of larger mass thermonuclear reaction – requires temp of 40,000,000 K to sustain 1 g of fusion fuel = 20 tons of coal occurs naturally in stars

86 Fission vs. Fusion 235 U is limited danger of meltdown toxic waste thermal pollution fuel is abundant no danger of meltdown no toxic waste not yet sustainable FISSIONFISSION FUSIONFUSION

87 Nuclear Power Fission Reactors Cooling Tower

88 Nuclear Power Fission Reactors

89 Nuclear Power Fusion Reactors (not yet sustainable)

90 Nuclear Power Fusion Reactors (not yet sustainable) Tokamak Fusion Test Reactor Princeton University National Spherical Torus Experiment

91 Synthetic Elements Transuranium Elements –elements with atomic #s above 92 –synthetically produced in nuclear reactors and accelerators –most decay very rapidly

92 Radioactive Dating half-life measurements of radioactive elements are used to determine the age of an object decay rate indicates amount of radioactive material EX: 14 C - up to 40,000 years 238 U and 40 K - over 300,000 years

93 Nuclear Medicine Radioisotope Tracers –absorbed by specific organs and used to diagnose diseases Radiation Treatment –larger doses are used to kill cancerous cells in targeted organs –internal or external radiation source Radiation treatment using  -rays from cobalt-60.

94 Nuclear Weapons Atomic Bomb –chemical explosion is used to form a critical mass of 235 U or 239 Pu –fission develops into an uncontrolled chain reaction Hydrogen Bomb –chemical explosion  fission  fusion –fusion increases the fission rate –more powerful than the atomic bomb

95 Others Food Irradiation –  radiation is used to kill bacteria Radioactive Tracers –explore chemical pathways –trace water flow –study plant growth, photosynthesis Consumer Products –ionizing smoke detectors Am


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