H/Chemistry Atomic Theory

H/Chemistry Atomic Theory

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 à mole

Inventions

Which of the following were invented

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.

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

à 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.

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

Isotopes

1. What do the following symbols represent?

a. e- ______

b. n0 ______

c. p+ ______

2. Which subatomic particles are found in an atom’s nucleus? ______

3. Which subatomic particle identifies an atom as that of a particular element? ______

4. Explain why atoms are neutral even though they contain charged particles.

______

5. What do the numbers, 39, 40, and 41 after the element name potassium refer to?

______

6. Write the symbolic notation for each of the following isotopes:

a. potassium-39 ______

b. potassium-40 ______

c. potassium-41 ______

7. Write an equation showing the relationship between an atom’s atomic number and its mass

number.

______

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.

Practice with Average Atomic Mass - Marbles

I have 3 marbles, weighing 1.59 g, 1.51 g, and 1.76 g.

To find the average mass of the marbles,

I add the masses up, then divide by 3, = 1.62 g.

The same calculation can be done as

4.86 g x 0.3333… = 1.62 g.

What about if I had 100 marbles, 79 marbles with a mass of 1.59 g, 10 with a mass of 1.51 g and 11 with a mass of 1.76 g? Which is the average mass?

79 x 1.59 g = 126 g

10 x 1.51 g = 15.1 g

11 x 1.76 g = 19.4 g

Total mass = 160.5 g divided by 100 = 1.60 g

Practice with Average Atomic Mass - Atoms

What is the average atomic mass of antimony? The isotopes of antimony and their percent abundance are Sb-121 (120.90 amu, 57.21%) and Sb-123 (122.90 amu, 42.79%)

What is the average atomic mass of vanadium? The isotopes of vanadium and their percent abundance are V-50 (49.95 amu, 0.250%) and V-51 (50.94 amu, 99.750%).

Atomic #, Mass # and Average Atomic Mass

IV. Nuclear Chemistry and Radioactivity

Nuclear reaction à 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 à 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”

Types of radiation

1. Alpha radiation (most common in elements with atomic # > 83)

Alpha particles = 2 p+ + 2 n0 (He nucleus, 42He, a) with 2+ charge

e.g. 22688Ra à 22286Rn + 42He (+ energy)

2. Beta radiation (most common in elements with high n0:p+ ratio)

Beta particles = 1 e- (0-1b) with 1- charge

Neutron à proton + beta particle

10n à 11p + 0-1b

e.g. 146C à 147N + 0-1b (+ 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 (00g)

-  usually accompany alpha and beta radiation

e.g. 23892U à 23490Th + 42He + 2 00g

Nuclei with lower neutron:proton ratios than optimal:

4. Positron Emission à 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 à neutron + positron

11p à 10n + 01b

e.g. 116C à 115B + 01b

5. Electron Capture

(most common in elements with a high n0:p+ ratio)

à more neutrons by pulling in an e- which combines with a proton to form a neutron

Proton + electron à neutron

11p + 0-1e à 10n

e.g. 0-1e + 8137Rb à 8136Kr + X-ray photon

Radioactive Particles WS

Positron same mass as e-’s 0+1b 1/1840 1+

Electron capture electrons 0-1b 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.

______

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: 23592U + 10n à 23692U à 9236Kr + 14156Ba + 3 10n + energy

(unstable)

The new neutrons (10n) à 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 à 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 à He + energy

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

Other fusion reactions occur in the sun:

He + He à Be + g (gamma ray)

He + Be à C + g

2. Synthetic Elements

a. Nuclear bullets

i. Bombard nuclei of elements with small particles such as p+, n , He (a particles)

& e- (0-1b particles)

ii.  Used to create Elements # 93-100

iii.  1919 first experiment: N + He à O + H (Rutherford)

b. Crashing nuclei

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

e.g. C + Cm à 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 24494Pu + 4820Ca à 289114Fl + 3 10n (1999, Russia)

Production of the Transuranium Elements

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:

21483Bi à 42He + _____

23993Np à 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.

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 à 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 à electromagnetic force

2. Speed of EMR is always the same: c = 3.00 x 108 m/s ß MEMORIZE

3.  e.g. light, microwaves, TV, radio, X-rays

B.  Remember:

Waves have wavelength, l (lambda) in nm and frequency, n (nu) or f

# waves/sec = Hz (hertz) = cycle/sec or s-1

** c = ln ß MEMORIZE

note: l and n are inversely proportional.

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

c = l n

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)

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.