235U P = atomic number = 92

N = Mass Number - Atomic Number = 235 - 92 = 143

e^- = Electrically neutral, ˆ number of electrons = number of protons

3-2 What is the basic structure of an atom?

1. Infer the existence of atoms from the laws of definite composition, conservation of mass, and multiple proportions.

2. List the five basic principles of Dalton’s atomic theory.

3. Describe models of the atom.

4. Compare and contrast the properties of electrons, protons, and neutrons.

5. Explain the particle-wave nature of electrons.

6. Describe the quantum model of the atom.

A. Creating Atomic Models By Inference: Relate to the Scientific Method: Uncertainty Principle also Inference.

1. Straw men: Develop a model and then try to destroy it.

2. Wind: Can’t be seen, but its force can be felt. The evidence of the wind is indisputable. This is inference.

3. Study the patterns of nature.

4. Develop models that fit the information.

5. Test the models.

B. Known Laws of Nature.

1. Law of definite composition: a compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. Transparency: The Law of Definite Composition
Example: Table sugar (sucrose) is composed of 42.1% carbon, 51.3% oxygen, and 6.5% hydrogen. The proportions are always the same.

2. Law of conservation of mass: In a chemical reaction, the mass of the reactants is equal to the mass of the products (mass is neither created nor destroyed). Transparency 3-10: Law of conservation of mass

3. Law of multiple proportions: the mass ratio for one of the elements that combines with a fixed mass of the other element can be expressed in small whole numbers. Transparency 3.11 Law of multiple proportions.

4. Taken together these three laws represent much of the quantitative data obtained by chemists in the 1700’s.

C. English chemist, John Dalton, argued that the three experimental laws could not be explained without assuming that all compounds are made from tiny particles such as atoms. This reasoning led to the development of the atomic theory, published in the early 1800’s.

While some exceptions to Dalton’s atomic theory were eventually discovered, the theory itself has never been discarded, only modified and expanded as the world of the atom was explored.

A. Testing the theory.

1. Scientists developed a tool called a cathode ray tube. Basically they shot electrons across a gas at low pressure. What they observed was a light. Transparency

a.

Cathode: negative electrode (-)

b. Anode: positive electrode (+)

c. Cathode Ray: light between the Cathode and Anode. This ray always originated at the cathode and traveled to the anode.

2. Cathode rays are what "paint" the pictures on television.

3.

As scientists studied the cathode ray tube they discovered that the atom is not indivisible after all. Instead it consists of smaller particles.

B. In 1897, the English physicist J. J. Thomson discovered that electrically charged plates and magnets deflected the straight paths of cathode rays. The direction of deflection shows that the particles making up cathode rays must be negatively charged. Transparency

C. The English physicists, G Johnstone Stoney named the small, negatively charged particles discovered in the cathode ray tube experiments electrons.

1. Later experiments determined the mass and charge of an electron. The electron was discovered to have a mass of nearly 2000 times smaller than that of hydrogen.

D. Atoms were known to be electrically neutral. This meant that an atom must contain some positively charged matter to balance the negative charges of its electrons.

J. J. Thomson developed a model of the atom that was based upon all of this information. It was called the plum pudding model. Thomson envisioned the atom as a ball of positive charge with negatively charged electrons embedded inside.

F. A student of J. J. Thomson, Ernest Rutherford, assembled a research team to perform an experiment that ultimately disproved the plum pudding model of the atom.

1. The name of the experiment is the Gold Foil Experiment. Transparency 3-16: Gold Foil Experiment Rutherford’s team directed a beam of tiny positively charged particles, called alpha particles, at a very thin gold foil sheet. The gold sheet was hammered extremely thin so that it would have relatively few atoms in thickness. The idea was to test if the atom was solid or not.

2. The team found that most of the alpha rays went directly through the gold foil. A few rays were deflected from their straight-line paths. But what really surprise the research team was that some of the alpha particles were deflected straight back!

3. Rutherford reasoned that the deflections resulted from electrical repulsion between the positive alpha particle and the positive charged matter contained in the atom.

4. This disproved the plum pudding model. If the positive charge were spread out within the atom, as in the plum pudding model, the backward scattering of alpha particles would not have been possible.

5. Because most of the alpha particles went through the gold foil Rutherford reasoned that atoms are composed mostly of space, with a small, dense, positively charged core. Figure 3-17 pg 86

6. This tiny center was named the nucleus, from the Latin word meaning "little nut".

a. Most of the mass is in the nucleus.

1) Protons (1 amu) are in the nucleus.

2) Neutrons (1 amu) are in the nucleus.

b. The nucleus is tiny compared to the volume of the atom. If the nucleus were the size of a marble, then the whole atom would be the size of a football stadium. Figure 3-18

G. Rutherford developed the Planetary Model of the atom. He supposed that electrons traveled in the space surrounding the nucleus in a way similar to the motion of the planets around the sun.

 

1. A maximum number of 7 Primary Shells.

2. The difference between these Primary Shells is the distance from the nucleus. The farther from the nucleus the electron is the more energy the electron has. Transparency: Rutherford's Planetary Model

H. In 1913, a young Danish physicist named Niels Bohr proposed that electrons could reside only in certain energy levels. Bohr transformed the way we think about electrons and modified the planetary model. The Quantum Theory Handout

1. An analogy that is often used is the ladder. Just like you cannot climb up the ladder in between the rungs, so electrons cannot exist between energy levels. On the side of pg 87

2. Bohr reasoned this from studying the bright-line emission spectrum of the elements, particularly hydrogen.

3. Running high voltage current through a gaseous form of the element creates the bright-line emission spectrum of elements.

I. More About Electrons: The Physics of Energy. In order to understand Bohr's thoughts

1. Bohr studied the bright line spectrum of elements.

a. He noticed that element has every a unique bright line spectrum.

b. He decided to focus on Hydrogen.

2. Before we go into his study of hydrogen, it is important to understand some simple light and wave physics.

a. Waves: When we describe a wave we usually describe the following features of a wave. These are generally true of all waves, but particularly true of light.

1) Wavelength (λ): The distance between two identical portions of a wave.

2) Frequency (f): The number of times a wave passes a particular spot in a set period of time. The unit for this is hertz (Hz = 1/s)

3) Axis: The midline of a wave.

4) Peak: The top of the wave.

5) Trough: The bottom of the wave

6) Amplitude: The measurement from axis of a wave to either the peak or the trough.

 

 

 

 

 

b. Relationships between these features:

Frequency and Wavelength were found to be inversely proportional. \ f » 1/l

2) The constant for this equation turned out to be the velocity (v) of the wave therefore: f = v/l and when looking at light it is: f = c/l where c = 3.00E8 m/s (speed of light).

3) Energy was found to be directly proportional to frequency: \ E » f.

4) Since they are proportional they can be made equal by multiplying by a constant. A man named Plank discovered the value of this constant so it is called Plank's Constant (h). E = hf.

5) Since most of the time we compare energy and wavelength, by substituting in wavelength for frequency we get: E=hc/l .

c. Bright Line Spectrum:

1) Language:

Excited State: Electron has absorbed energy and gone to a higher energy level.

b) Ground State: Electron is in its normal state of energy.

c) Quantum: A specific amount of energy.

d) In nature things tend to go to the lowest energy level possible.

e) This means that the electrons would stay at the higher energy level for a moment and then go back to the lower energy level.

f) We usually describe this in terms of stability: Low energy = high stability, high energy = low stability.

2) The electron in ground state would absorb energy and go to an excited state where it would be unstable. When the electron returns to ground state it gives up the exact amount of energy it absorbed when becoming excited. This energy is given off in the form of a bright line of light.

a) Electron absorbs energy.

b) Electron goes to higher energy level.

c) Electron is unstable and goes to a lower energy level.

d) Electron gives off the exact amount of energy absorbed when it goes to a lower energy level.

e) Since there are specific wavelengths of light, electrons must only be able to absorb specific amounts (quantum) of energy!

J. Protons discovered

1. The positively charged nuclear particles that repelled the alpha particles in Rutherford's experiment were found to be very heavy. 2000X the mass of an electron.

2. Scientists called these protons. The mass of a proton presented a dilemma because the masses of all atoms besides hydrogen were known to be larger than the total mass of their protons and electrons. Clearly there must be a third particle.

K. Neutrons discovered

1. This discovery was a little more difficult and took over 30 years.

2. A British scientist, James Chadwick, found a penetrating beam that was made of particles that had approximately the mass of protons. Also, the beam was not deflected by electric or magnetic fields. Chadwick deduced that the beam was composed of neutrons—neutral particles that have mass equal to that of protons.

3-2. How do the structures of atoms differ?

A. Electron Configuration PowerPoint Presentation