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Unit 3: “Atomic Structure”

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

2 I. Atomic Structure A. Section 4.1 Defining the Atom
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

3 1. Dalton’s Atomic Theory (experiment based!)
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. atoms identical different John Dalton (1766 – 1844) 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. compounds never

4 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

5 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

6 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

7 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

8 Modern Cathode Ray Tubes
Television Computer Monitor Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

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

10 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

11 Eugen Goldstein 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 proton James Chadwick neutron no charge equal

12 4. Subatomic Particles Particle Charge Mass (g) Location Electron (e-)
-1 9.11 x 10-28 Electron cloud Proton (p+) +1 1.67 x 10-24 Nucleus Neutron (no)

13 Electron cloud Nucleus

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

15 6. Ernest Rutherford’s Gold Foil Experiment - 1911
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

16 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 #2 Target #1

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

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

19 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

20 1. Atomic Number C. Section 4.3: Distinguishing Among Atoms identical
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. ___________________________ PROTONS atomic number number of protons # protons in an atom = # electrons

21 2. Definition: Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element # of protons Atomic # (Z) Carbon Phosphorus Gold 6 6 15 15 79 79

22 3. Mass Number Definition: Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p+ + n0 Nuclide p+ n0 e- Mass # Oxygen - 10 - 33 42 - 31 15 18 8 8 18 75 Arsenic 75 33 Phosphorus 16 15 31

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

24 Br 80 5. Symbols 35 a. Find each of these: 1) number of protons
2) number of neutrons 3) number of electrons 4) Atomic number 5) Mass Number 35 45 35 35 80

25 b. If an element has an atomic number of 34 and a mass number of 78, what is the:
number of protons number of neutrons number of electrons complete symbol 34 44 34 Se 78 34

26 Pt d. If an element has 78 electrons and 117 neutrons what is the
Atomic number Mass number 3) number of protons 4) complete symbol 78 195 78 Pt 195 78

27 6. Isotopes 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

28 Isotopes Frederick Soddy
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. Isotopes same element

29 a. We can also put the mass number after the name of the element:
7. Naming Isotopes a. We can also put the mass number after the name of the element: b. Examples: carbon-12 carbon-14 uranium-235

30 same element Isotopes different
c. _______ are atoms of the ___________ having ________ masses, due to varying numbers of neutrons. different Isotope Protons Electrons Neutrons Nucleus Hydrogen–1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2

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

32 9. IONS 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

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

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

35 Learning Check – Counting
State the number of protons, neutrons, and electrons in each of these ions. 39 K+ 16O Ca +2 #p+ ______ ______ _______ #no ______ ______ _______ #e- ______ ______ _______ 19 8 20 20 8 21 18 10 18

36 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- ___________

37 Charges on Common Ions -3 -2 -1 +1 +2
By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.

38 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

39 10. Atomic Mass 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

40 11. Measuring Atomic Mass 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

41 12. To calculate the average:
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

42 Composition of the nucleus
13. Atomic Masses Atomic mass is the average of all the naturally occurring isotopes of that element. Isotope Symbol Composition of the nucleus % in nature Carbon-12 12C 6 protons 6 neutrons 98.89% Carbon-13 13C 7 neutrons 1.11% Carbon-14 14C 8 neutrons <0.01% Carbon = 12.01

43 - Page 117 Question Knowns and Unknown Solution Answer

44 . 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

45 15. The Periodic Table: A Preview
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

46 c. Each vertical column is called a ____________
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

47 A. Section 6.1: Organizing the Elements NOTES:
II. The Periodic Table A. Section 6.1: Organizing the Elements NOTES: 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 1700. As more were discovered, chemists realized they needed a way to ________ the elements. 13 organize

48 properties 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 triads

49 1. Mendeleev’s Periodic Table
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

50 2. Mendeleev He left blanks
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

51 3. A Better Arrangement Moseley
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 atomic number


53 4. The Periodic Law 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

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

55 Areas of the Periodic Table
gases 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? _________: border the line-2 sides Properties are __________ between metals and nonmetals Metalloids intermediate

56 Squares in the Periodic Table
symbols 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) names number mass blue = liquid


58 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 reacting alkaline earth metals halogens noble gases

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!

60 Chemistry – Unit 4 Chapter 25 Nuclear Chemistry 60

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

62 Nuclear Binding Energy
Energy released when a nucleus is formed from nucleons. High binding energy = stable nucleus. E = mc2 E: energy (J) m: mass defect (kg) c: speed of light (3.00×108 m/s)

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

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

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

66 Nuclear Decay Beta Emission electron Positron Emission positron

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

68 IQ# 1 Balance the following equations:

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

70 Belt of Stability and Radioactive Decay

71 Half-life Half-life (t½)
Time required for half the atoms of a radioactive nuclide to decay. Shorter half-life = less stable.

72 Half-life mf: final mass mi: initial mass n: # of half-lives

73 Half-life t½ = 5.0 s mi = 25 g mf = ? total time = 60.0 s
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: t½ = 5.0 s mi = 25 g mf = ? total time = 60.0 s n = 60.0s ÷ 5.0s =12 WORK: mf = mi (½)n mf = (25 g)(0.5)12 mf = g

74 Example: How much of a 500. g sample of Uranium-235 would be left after five half-lives?
Mi = 500 g n = 5 (n = # of half-lives) Mf = ? mf = mi (½)n mf = (500 g)(0.5)5 mf = 15.6 g

75 Example: A 16. 00 mg sample of Radon-222 decays to 0
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

76 mf = mi (½)n mf = (1.000 mg)(0.5)5 mf = 0.03125 mg
Example: The half-life of molybdenum-99 is 67 hours. How much of a mg sample is left after 335 hours? Mi = mg mf = mi (½)n mf = (1.000 mg)(0.5)5 mf = mg Half-life = 67 h Rxn time = 335 h Mf = ? n = 335 / 67 = 5

77 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? mf = mi (½)n mf = (64 mg)(0.5)3 mf = 8.0 mg Mi = 64 mg n = 3 Mf = ?

78 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. Enter “50” in Shake # 0 row for Trial 1, 2, & 3 and “150” for  (Sum of) of trials. 4. 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.

79 Prepare a graph in your lab book:
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.

80 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)

81 · · · · Half-Life Lab Title At least 2/3 of pg Use Ruler for axis
150 Title At least 2/3 of pg Use Ruler for axis Label Axis Circle Data Points Best Fit Curve 75 Σ of Trials Convenient #’s 1 + 1 + 1.2 = 3.2 ~ 1.1 shake 3 37.5 18.75 Half-life = 1.1 shake 1 2 3 4 5 6 7 8 9 10 1 2 3 Shake #

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

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

84 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

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

86 Nuclear Power Cooling Tower Fission Reactors

87 Nuclear Power Fission Reactors

88 Nuclear Power Fusion Reactors (not yet sustainable)

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

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

91 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: 14C - up to 40,000 years 238U and 40K - over 300,000 years

92 Radiation treatment using
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.

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

94 Others Food Irradiation Radioactive Tracers Consumer Products
 radiation is used to kill bacteria Radioactive Tracers explore chemical pathways trace water flow study plant growth, photosynthesis Consumer Products ionizing smoke detectors - 241Am

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