I. Historical Backgroundtamchemistryhart.weebly.com/uploads/3/8/0/0/38007377/atomic_the… · • How are light and electrons related? • How do we “see” where electrons are

H/Chemistry Atomic Theory

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Atomic Theory Essential Questions:

• What are we made of?

• How are scientific models developed?

• Do atoms exist or are they just concepts invented by scientists? What evidence is there in

your everyday life for the existence of atoms?

• How did the understanding of the atom affect historical events?

• How have historical events affected the model of the atom?

• What do we think the atom “looks like” now?

• If the atom is mostly empty space, why doesn’t my butt fall through the chair?

• How are light and electrons related?

• How do we “see” where electrons are located in the atom?

• Why is the location of electrons so important?

I. Historical Background A. The Greek philosophers

Democritus (460-370 BC) – “atomism”

Aristotle (384-322 BC) and other philosophers –earth, air, fire and water, denied the

existence of atoms

No experiments

B. On the way to chemistry – alchemy and phlogiston Alchemy – transmutation of other metals into gold (~ Middle Ages and before)

Phlogiston – an imaginary element, believed to separate from every combustible

body in burning (1700s)

C. The first experimental chemists – focus on gases Henry Cavendish (1731-1810) – discovered hydrogen

Joseph Priestley (1733-1804) - oxygen

Antoine Lavoisier (1743-1794) - oxygen

Karl Wilhelm Scheele (1742-1786) - oxygen

Count Amedeo Avogadro (1776-1856) - gases mole


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Inventions Which of the following were invented

a) before 1800?

b) between 1800 and 1900?

c) after 1900?

D. Three Laws

1. Law of Conservation of Mass Antoine Lavoisier (1770s) Mass of reactants = mass of products (relates to before and after reaction)

2. Law of Definite Proportions Joseph Proust (1799) Compounds always contain elements in the same proportion by mass, no matter how they are made or where they are found.

3. Law of Multiple Proportions John Dalton (1803) Different compounds made from the same elements:

The ratio of mass of an element in the first compound relative to the mass of the same element in a second compound is a fixed whole number.


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E. Dalton’s Atomic Theory

1. All matter is made of extremely small particles called atoms. 2. All atoms of a given element are identical (mass, physical and chemical

properties).

3. Atoms of different elements have different masses, and different physical and chemical properties.

4. Different atoms combine in simple whole number ratios to form compounds.

5. In a chemical reaction, atoms are combined, separated, or rearranged in

chemical reactions.

6. Atoms cannot be created, divided into smaller particles or destroyed.

II. Subatomic Particle Discovery

A. Electron was the first discovered!

1897 – J.J. Thomson 1. Experiment: cathode rays

a. passed electricity through a partially-evacuated tube of gas b. observed a ray of light crossing from one electrode to the other c. the ray of light moved a paddle wheel inside the tube

2. Conclusions:

a. Since ray moved toward (+) electrode, the ray must be (-) charged. b. Since the paddle wheel moved, the ray must be made up of particles. c. Since Thomson performed the experiment with different gases, with the

same results, the particles must be in all atoms.

B. Nucleons

1. Protons: 1909/1910 – Rutherford (specialty = radioactivity) (Geiger and Marsden performed the experiments; Rutherford interpreted them) Alpha particles (He nucleus, + charge) shot through gold foil were deflected in peculiar ways, inconsistent with the plum pudding model of the atom area of concentration of positive charge

2. Neutrons: 1932 – James Chadwick

(worked with Hans Geiger, then Ernest Rutherford) Uncharged particles in the nucleus with mass were pushed out of beryllium when bombarded with alpha particles. These particles accounted for the “missing mass” in the nucleus.


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III. Subatomic Particles

A.Electrons

B.Nucleons

(1) Protons

(2) Neutrons

Atomic Mass

Atomic number = # protons (also = # electrons)

Mass number = # protons + # neutrons

# neutrons = mass number – atomic number

Isotopes # protons # neutrons # electrons

Hydrogen-1 (protium) 11H 1 0 1

Hydrogen-2 (deuterium) 21H 1 1 1

Hydrogen-3 (tritium) 31H 1 2 1

POGIL – Isotopes

POGIL - Ions

Location Charge Relative mass

Symbol


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Average Atomic Mass 1 atomic mass unit (amu) = 1/12 mass of a C-12 atom Calculate average atomic mass by using weighted averages which take into account the relative abundance of each isotope. Note that what is on the periodic table is the average atomic mass, not the mass of a single isotope.

POGIL – Average Atomic Mass

Atomic #, Mass # and Average Atomic Mass

Atomic # Mass # Average Atomic Mass

# p+ # p+ + #n0 Weighted average of all isotopes

whole #

Whole number Decimal, limited sfs

found on PT

NOT on PT Found on PT

IV. Nuclear Chemistry and Radioactivity Nuclear reaction change in the identity of the element(s) Radioactivity = radiation emitted spontaneously by certain elements whose nuclei contain an unstable neutron:proton ratio Smaller elements Stable neutron:proton ratio = 1 n0:1 p+ for masses < 20 (i.e. mass # = 2 x atomic #) Largest elements Stable neutron:proton ratio = 1.5 n0:1 p+ for largest atoms All elements with atomic #s > 83 are radioactive Unstable nuclei emit radiation and change their identities (“radioactive decay”) Historical figures: Wilhelm Roentgen (1845-1923) - discovered X-rays –1895 Henri Becquerel (1852-1908) - discovered radioactivity in U - late 1800s Ernest Rutherford (1871-1937) – id’d different types of radiation (beg. 1898) Pierre (1867-1934) and Marie Curie (1859-1906) - discovered radium and polonium – 1898; first used the term “radioactivity”


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Types of radiation

1. Alpha radiation (most common in elements with atomic # > 83) Alpha particles = 2 p+ + 2 n0 (He nucleus, 4

2He, α) with 2+ charge e.g. 226

88Ra 22286Rn + 42He (+ energy)

2. Beta radiation (most common in elements with high n0:p+ ratio) Beta particles = 1 e- (0

-1β) with 1- charge Neutron proton + beta particle 1

0n 11p + 0-1β

e.g. 14

6C 147N + 0-1β (+ energy)

Note: The sum of the mass #s and atomic #s on both sides of the equation are the same. 3. Gamma radiation Gamma rays = high-energy radiation with no mass and no charge (0

0γ) - usually accompany alpha and beta radiation

e.g. 238

92U 23490Th + 42He + 2 00γ

Nuclei with lower neutron:proton ratios than optimal: 4. Positron Emission more neutrons by converting a proton into a neutron (most common in lighter elements with low n0:p+ ratio) Positron = particle with same mass as an e-, but opposite charge Proton neutron + positron 11p 10n + 01β e.g. 11

6C 115B + 01β

5. Electron Capture (most common in elements with a high n0:p+ ratio) more neutrons by pulling in an e- which combines with a proton to form a neutron Proton + electron neutron 1

1p + 0-1e 10n e.g. 0

-1e + 8137Rb 81

36Kr + X-ray photon


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Radioactive Particles WS

Positron same mass as e-’s 0+1β 1/1840 1+ Electron capture electrons 0-1β 1/1840 1-

(Added to the reactants side)

1. Which radioactive emission has the greatest mass? __________________________ Least mass? _______________________________ 2. Why do you think gamma rays are drawn as wavy lines? __________________________ _______________________________________________________________________ 3. To which charged plate are the alpha particles attracted? Explain. ______________________________________________________________________ 4. To which charged plate are the beta particles attracted? _______________________________________________________________________ Why do the beta particles have a greater curvature than the alpha particles? ______________________________________________________________________ 5. Explain why the gamma rays do not bend toward one of the electrically charged plates. ______________________________________________________________________


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Nuclear Fission and Fusion Nuclear fission = the splitting of a nucleus into smaller, more stable fragments, accompanied by a large release of energy e.g. Uranium-235: 235

92U + 10n 236

92U 9236Kr + 141

56Ba + 3 10n + energy (unstable) The new neutrons (1

0n) fission of more U-235 (= chain reaction, a self-sustaining process) Chain reaction requires a critical mass (= minimum amount of starting material to maintain a chain reaction); supercritical mass may violent nuclear explosion results in radioactive waste Practical examples = nuclear power plant, atomic bomb Nuclear Fusion = the process of binding smaller atomic nuclei into a single larger and more stable nucleus, requiring a huge amount of energy to initiate, followed by a large release of energy 1. Creation of natural elements

a. Hydrogen, other light elements - from the Big Bang

b. Elements #2-92 (except Fr, Pr, Te, At) Nuclear fusion occurs in stars (naturally) Occurs in hydrogen bomb (artificially) > 2 x 107oC The sun converts 3 x 1014 g of H into He every second.

4 H He + energy

Mass is not conserved. Mass is converted into energy via E = mc2

Other fusion reactions occur in the sun:

He + He Be + γ (gamma ray) He + Be C + γ

11

4 2

4 2

4 2

4 2

84

84

12 6


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2. Synthetic Elements

a. Nuclear bullets

i. Bombard nuclei of elements with small particles such as p+, n , He (α particles) & e- (0

-1β particles)

ii. Used to create Elements # 93-100

iii. 1919 first experiment: N + He O + H (Rutherford) b. Crashing nuclei

i. Accelerators hurl nuclei into each other at very high speeds.

e.g. C + Cm No + 2 n carbon curium nobelium neutron

ii. Elements beyond # 100

iii. These elements are very unstable:

e.g. Element 109 existed for only 3.4 x 10-3 sec (3 atoms)

c. Superheavy elements (“transuranium” elements)

Stability of nucleus of atom depends on filling "shells" within nucleus with alternating p+ and n. The more filled shells, the more stable it would be.

e.g. Element 114 244

94Pu + 4820Ca 289

114Fl + 3 10n (1999, Russia)

Production of the Transuranium Elements

12 6

244 96

254 102

14 7

4 2

17 8

11

1 0

4 2


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1. Does the diagram illustrate a natural transmutation reaction or an induced transmutation reaction? ______________________________________________________________________________ 2. What is the name and nuclear symbol of the isotope produced in the reaction? ____________________________________________________________ 3. Write nuclear equations to show how dubnium-263, lawrencium-262, and seaborgium-266 can be produced from a nuclear reaction of neon-22 and americium-244. _____________________________________________________________________________ _____________________________________________________________________________ _____________________________________________________________________________ 5. Each of the radioisotopes in the table decays within 20 seconds to 10 hours. Write a nuclear equation for each decay. _______________________________________________________________________________ _______________________________________________________________________________ ________________________________________________________________________________ ________________________________________________________________________________ 6. Which, if any, of the four isotopes listed in the table would you expect to find at Earth’s surface? Why? __________________________________________________________________________

Nuclear equations

Complete the following equations : 214

83Bi 42He + _____

239

93Np 23994Pu + ______

Write a balanced nuclear equation for the alpha decay of americium-241. Write a balanced nuclear equation for the beta decay of bromine-84.


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Next Steps in Discovering the Structure of Atoms – Properties of Electrons

• Wave nature of light – EMR (James Maxwell, 1864)

• Particle nature of light – quantum (Max Planck, late 1800s)

• Emission of light and other EMR from heated elements emission spectra V. Electromagnetic Radiation (EMR)

A. Definition: energy that exhibits wave-like (or oscillating) behavior as it travels through space 1. James Maxwell in 1864 – unified the electric and magnetic forces

electromagnetic force 2. Speed of EMR is always the same: c = 3.00 x 108 m/s MEMORIZE 3. e.g. light, microwaves, TV, radio, X-rays

B. Remember:

Waves have wavelength, λ (lambda) in nm and frequency, ν (nu) or f # waves/sec = Hz (hertz) = cycle/sec or s-1 ** c = λν MEMORIZE note: λ and ν are inversely proportional.

Electromagnetic Radiation Practice Problems from Textbook (pp. 121, 124)

c = λ ν Microwaves are used to transmit information. What is the wavelength of a microwave having a frequency of 3.44 x 109 Hz? (8.72 x 10-2 m)

1. What is the frequency of green light which has a wavelength of 4.90 x 10-7 m?

(6.12 x 1014 s-1)


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2. An X-ray has a wavelength of 1.14 x 10-10 m. What is its frequency? (2.63 x 1018 s-1) 3. What is the speed of an electromagnetic wave that has a frequency of 7.8 x 106 Hz?

(3.00 x 108 m/s) 4. A popular radio station broadcasts with a frequency of 94.7 MHz.

What is the wavelength of the broadcast? (1 MHz = 106 Hz) (3.17 m) The Idea of the Quantum

• Quantum = the smallest discrete amount of energy that can exist independently, esp. as EMR • 1 quantum of light = 1 photon E = hν h = Planck’s constant = 6.626 x 10-34 J.s

(Energy emitted = Planck’s constant x frequency of light hitting it)

∴ there is a direct relationship between E and ν

• The amount of energy in EMR increases with increasing frequency:

E = hν Planck’s constant = h = 6.626 x 10-34 J·s

Example: Tiny water drops in the air disperse the white light of the sun into a rainbow. What is the energy of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 1014 s-1? (4.79 x 10-19 J) 5. What is the energy of each of the following types of radiation? a. 6.32 x 1020 s-1 (4.19 x 10-13 J)

b. 9.50 x 1013 Hz (6.29 x 10-20 J) c. 1.05 x 1016 s-1 (6.96 x 10-18 J)

6. Name the types of radiation in each part of #5.


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Electromagnetic Radiation Practice Problems 1. What is the frequency of EMR with a wavelength of 235 pm? What type of EMR is this? 2. What is the frequency of EMR with a wavelength of 0.614 cm? What type of EMR is this? 3. What is the wavelength of EMR with a frequency of 8,512 Hz? What type of EMR is this? 4. What is the wavelength of EMR with a frequency of 625 x 1017 Hz? What type of EMR is this? 5. If the speed of light is 3.00 x 108 m/s, calculate the wavelength of the electromagnetic radiation whose

frequency is 7.500 x 1012 Hz. 6. Determine the frequency of light with a wavelength of 4.257 x 10-7 cm. 7. For the following sources, calculate the missing member of the wavelength/frequency pair. a) FM radio waves with a frequency of 94.7 Hz. b) A laser with a wavelength of 1064 nm. c) An X-ray source, emitting X-rays with a wavelength of 175.4 pm.


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VI. Bohr’s model Neils Bohr (Danish physicist)

– studied with both Thomson and Rutherford – refined Rutherford’s model and published his own model in 1913 – incorporated Planck’s idea of quanta of energy – received the Nobel Prize in 1922 – provided an explanation for the spectral lines of hydrogen

POGIL – Electron Energy and Light 1. Electrons are arranged in circular paths, “orbits” around the nucleus. 2. Electrons have fixed energy levels which prevent them from falling into the nucleus.

Electrons closest to the nucleus have lowest Etotal = KE + PE (most stable) 3. To move from one energy level to another, electrons must gain or lose energy.

EMR is emitted from the atom when electrons fall down to a lower energy level. 4. The distance between energy levels is not equal.

A “quantum” is the amount of energy needed to make the leap. 5. Bohr’s model did not explain the spectra of atoms with > 1 electron.

Definitions related to Spectra

Spectrum = whole range of related qualities [Latin: appearance, from specere – to view] Electromagnetic spectrum = all EMR arranged according to λ Emission = any radiation of energy by EM waves [Latin: emitto – to send out, to utter, to hurl, to set free] Emission spectrum = the spectrum into which light or other EMR from any source can be separated Continuous spectrum = a spectrum whose source emits light of every λ in a continuous band

Bright-line spectrum = pattern of bright lines on a dark background. Source = glowing gas that radiates in special λ’s characteristic of the chemical composition of the gas

Line spectrum

_______E2________

_______E1________

Energy given off each photon (quantum of light) emitted has a characteristic λ which contributes a line to the spectrum. Each substance (element) has a set of λ’s as distinct as a fingerprint.

Visible Balmer Series e- drop to n=2


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VII. Quantum Theory of the Atom (current model) A. Quantum Mechanics

Louis de Broglie: discovered the wave nature of electrons Erwin Schrödinger, 1926: complex mathematical equation which treats electrons as waves foundation of quantum theory 1. Uses probabilities (90%) 2. Heisenberg Uncertainty Principle: we cannot simultaneously measure an e-‘s

position and its velocity B. Orbitals and Quantum #’s

1. All probable locations of an electron = orbital in the shape of an “electron cloud” 2. Each e- in an atom has its own set of four quantum #’s: n, l, m, s 3. Principal energy levels (n) correspond to Bohr’s orbits. Total # e- in one principal energy level = 2n2, where n = principal energy level

4. Sublevels (l), “magnetic” (position in space: m) and spin (s) are additions and take

into account the observed energies of the e-s.

Orbitals

1. What is the shape of an s orbital? ___________________________________________________

2. What is the relationship between the size of an s orbital and the principal energy level in which it is found? ______________________________________________________________________________


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3. What is the shape of a p orbital? How many p orbitals are there in a sublevel? ______________________________________________________________________________ 4. How many electrons can each orbital hold? _____________________ 5. Look at the diagrams of the p orbitals. What do x, y, and z refer to? ______________________________________________________________________________ 6. How many d orbitals are there in a given sublevel? How many total electrons can the d orbitals in a sublevel hold? ______________________________________________________________________________ 7. Which d orbitals have the same shape? _____________________________________________ 8. What point in each diagram represents an atom’s nucleus? ______________________________ 9. How likely is it that an electron occupying a p or a d orb ital would be found very near an atom’s nucleus? What part of the diagram supports your conclusion? _____________________________________________________________________________ _____________________________________________________________________________

Principal Energy Level (n)

(or “Shell”)

Sublevel (l) (or “Subshell”)

shape of probability cloud

3-D (m) (position in

space) x, y, z

Spin (s) + ½ - ½

Total # e-

2n2

1

s only 1 orientation

1 e- + ½ 1 e- - ½

2(1)2 = 2

2

s, p 3 orientations for p: px, py, pz

2(2)2 = 8

3

s, p, d etc. 2(3)2 = 18

4

s, p, d, f 2(4)2 = 32

5

s, p, d, f 2(5)2 = 50

6

s, p, d, (f) 2(6)2 = 72

Sublevels (# orbitals in each sublevel)

A. s: holds 2 e- x 1 orbital = 2 e- total B. p: holds 2 e- x 3 orbitals = 6 e- total C. d: holds 2 e- x 5 orbitals = 10 e- total D. f: ? holds 2 e- x 7 orbitals = 14 e- total


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VIII. Electron Configuration – how electrons are distributed within an atom Three Rules for Filling Orbitals

1. Aufbau: Fill in order of increasing energy levels.

2. Hund’s Rule: Fill all orbitals at same energy level with at least 1 e-, before adding the

second e-.

3. Pauli Exclusion Principle: only 2 e- per orbital, of opposite spin

1. What does each small box in the diagram represent? _____________________ 2. How many electrons can each orbital hold? _________________ 3. How many electrons can the d sublevel hold? __________________ 4. Which is associated with more energy: a 2s or a 2p orbital? ______________ 5. Which is associated with more energy: a 2s or a 3s orbital? ______________ 6. According to the Aufbau Principle, which orbital should fill first, a 4s or a 3d orbital? _________ 7. Which orbital has the least amount of energy? __________________ 8. What is the likelihood that an atom contains a 1s orbital? ___________________________ 9. Sequence the following orbitals in the order that they should fill up according to the Aufbau Principle: 4d, 4p, 4f, 5s, 6s, 3d, 4s. _____________________________________________________________________________

Documents

I. Historical Backgroundtamchemistryhart.weebly.com/uploads/3/8/0/0/38007377/atomic_the… · • How are light and electrons related? • How do we “see” where electrons are