1. Chapter 4 Models of the Atom by Christopher G. Hamaker
Illinois State University
© 2014 Pearson Education, Inc. 1

2. Dalton Model of the Atom
Greeks first considered that matter is composed of
small indivisible particles and called them atomos (indivisible)
The theory reemerged again in the early nineteenth
century, championed by John Dalton.
John Dalton proposed that all matter is made up of tiny particles.
These particles are molecules or atoms.
Molecules can be broken down into atoms by chemical processes.
Atoms cannot be broken down by chemical or physical processes.
Dalton was an English chemist, physicist, and meteorologist.

3. Chemistry Connection: John Dalton
Dalton began teaching at the age of 12 and developed an interest in science.
He was especially interested in meteorology and kept a lifelong daily journal of atmospheric conditions.
Although he was color blind, Dalton was able to make many great contributions to chemistry, including his model of the atom and the law of definite proportions.

4. Dalton Atomic Theory A Summary of Dalton Atomic Theory:
An element is composed of tiny, indivisible, indestructible particles called atoms.
All atoms of an element are identical and have the same properties.
Atoms of different elements combine to form compounds. Atoms are neither created nor destroyed in chemical reactions. It is called law of conservation of mass.
Eg. 8 H O2 → 8 H2O
4. Compounds contain atoms in small whole number ratios.
It is called law of definite proportion. A given compound
always contains exactly the same proportion of elements.
Eg. H2O, CO2, CH4
5. Atoms can combine in more than one ratio for forming different
compounds. It is called law of multiple proportions.
Eg. CO. CO2 & NO, NO2

5. Dalton Atomic Theory, Continued
The first two parts of atomic theory were later proven incorrect. We will see this later.
Proposals 3, 4, and 5 are still accepted today.
The Dalton theory was an important step in the further development of atomic theory.

6. Thomson Model of the Atom
Toward the end of the 1800s, evidence was seen that atoms were divisible.
Two subatomic particles were discovered.
Negatively charged electrons, e–.
Positively charged protons, p+.
An electron has a relative charge of –1, and a proton has a relative charge of +1.
Joseph John Thomson OM PRS was an English physicist and Nobel laureate in physics, credited with the discovery and identification of the electron; and with the discovery of the first subatomic particle

7. Thomson Model of the Atom, Continued
J. J. Thomson proposed a subatomic model of the atom in 1903.
Thomson proposed that the electrons were distributed evenly throughout a homogeneous sphere of positive charge.
This was called the plum pudding model of the atom.

8. Thomson Model of the Atom, Continued
Originally, Thomson could only calculate the charge-to-mass ratio of a proton (9.58 x 107 C/kg) and an electron (1.76  1011 C/kg).
Robert Millikan determined the charge of an electron (1.60 x C) in 1911.
Then Thomson calculated the masses of a proton and electron:
A proton has a mass of 1.67 × 10–24 g.
An electron has a mass of 9.11 × 10–28 g.
charge/mass ratio of the proton: x 107 coulomb/kg, (B) charge/mass ratio of the electron: 1.76  1011 coulombs/kg.
(C) the charge on an electron is x coulomb

9. Radioactivity
Radioactivity is the spontaneous emission of high-energy radiation by an atom of heavier elements.
Its discovery showed that the atom had more subatomic particles and energy associated with it.
Three types of radiation were discovered by Ernest Rutherford:
 particles (positively charged, α2+ or He2+)
 particles (negatively charged, like electrons, e− or β−)
 rays (uncharged, like photons)
Paul Villard, a French Chemist discovered the gamma rays.

10. equations for alpha decay
Pb = plumbum (Lead)
equations for beta decay
Pr = Praseodymium
Praseodymium
Gamma Radiation/Decay
Gamma rays is electromagnetic radiation similar to light. Gamma decay does not change the mass or charge of the atom from which it originates. Gamma is often emitted along with alpha or beta particle ejecton.

11. Rutherford Model of the Atom
Rutherford’s student fired alpha particles at thin gold foils. If the plum pudding model of the atom was correct, α particles should pass through undeflected.
However, some of the alpha particles were deflected backward.
Based on Thomson’s model, Rutherford expected that the positively charged alpha particles should pass throughout the uniform sphere of positively charged matter with little or no deflection.

12. Rutherford Model of the Atom, Continued
Most of the alpha particles passed through the foil because an atom is largely empty space.
At the center of an atom is the atomic nucleus, which contains the atom’s protons.
The alpha particles that bounced backward did so after striking the dense nucleus.
A) ɑ-particles are about 7300 time more massive than electrons.
B) If Thomson’s model was correct, maximum deflecting force on the
ɑ-particle as it passes near a positive charge will be far too small to
deflect the particle by even 1o.
C) Electrons in the atom would also have very little effect on the
massive, energetic ɑ-particle.

13. Rutherford Model of the Atom, Continued
Rutherford proposed a new model of the atom:
The negatively charged electrons are distributed around a positively charged nucleus.
An atom has a diameter of about 1 × 10–8 cm and the nucleus has a diameter of about 1 × 10–13 cm.
If an atom were the size of the Superdome, the nucleus would be the size of a marble.

14. Rutherford Model of the Atom, Continued
Based on the heaviness of the nucleus, Rutherford predicted that it must contain neutral particles in addition to protons.
Neutrons, n0, were discovered by James Chadwick. A neutron is about the size of a proton without any charge.

15. Atomic Notation
Each element has a specific number of protons in the nucleus. This is the atomic number, Z.
The total number of protons and neutrons in the nucleus of an atom is the mass number, A.
We use atomic notation to display the number of protons and neutrons in the nucleus of an atom:

16. Using Atomic Notation An example: Si
The element is silicon (symbol Si).
The atomic number is 14; silicon has 14 protons.
The mass number is 29; the atom of silicon has protons + neutrons = 29.
The number of neutrons is A – Z = 29 – 14 = neutrons. 29 14

17. Isotopes
All atoms of the same element have the same number of protons.
Most elements occur naturally with varying numbers of neutrons.
Atoms of the same element that have a different number of neutrons in the nucleus are called isotopes.
Isotopes have the same atomic number, but different mass numbers.

18. Isotopes, Continued The Isotopes of Hydrogen
We often refer to an isotope by stating the name of the element followed by the mass number.
The Isotopes of Hydrogen
How many protons and neutrons does an atom of lead-206 have? The atomic number of Pb is 82, so it has 82 protons. Pb-206 has 206 – 82 = 124 neutrons.

19. A certain isotope X contains 23 protons and 28 neutrons.
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
Number of protons = Number of electrons in a neutral atom
Number of neutrons = Mass number – Number of protons
EXERCISE!
A certain isotope X contains 23 protons and 28 neutrons.
What is the mass number of this isotope?
Identify the element
How many electrons are there?
Mass Number = = 51
The mass number is 51. Mass Number = # protons + # neutrons. Mass Number = = 51. The element is vanadium.

20. Simple and Weighted Averages
A simple average assumes the same number of each object.
A weighted average takes into account the fact that we do not have equal numbers of all the objects.
A weighted average is calculated by multiplying the percentage of the object (as a decimal number) by its mass for each object and adding the numbers together.

21. Average Atomic Mass
Since not all isotopes of an atom are present in equal proportions, we must use the weighted average.
Gallium has two isotopes:
69Ga, with a mass of amu and 60.11% abundance.
71Ga, with a mass of amu and 39.89% abundance.
The average atomic mass of Gallium is:
( amu)(0.6011) + ( amu)(0.3989) = amu

22. An element has two isotopes. First isotope is 62.60% with
EXERCISE!
An element has two isotopes. First isotope is 62.60% with
a mass of u and second isotope is 37.40% with a
mass of u.
Calculate the average atomic mass and identify the element.
Avg. At. Mass = (0.6260)( u) + (0.3740)( u) = u
The element is Rhenium (Re).
Avg. At. Mass = (0.6260)( u) + (0.3740)( u) = u

23. The Periodic Table
We can use the periodic table to obtain the atomic number and atomic mass of an element.
The periodic table shows the atomic number, symbol, and atomic mass for each element.

24. Chemistry Connection: Heavy Water
Heavy water still has the formula H2O, but the hydrogen atoms are the isotope hydrogen-2.
Hydrogen-2 is often referred to as deuterium, and is given the symbol D.
Heavy water (D2O) is slightly more dense than light water (H2O), and has slightly higher melting and boiling points.
Heavy water is used in nuclear reactors to slow down neutrons released during the fission process.

25. The Wave Nature of Light
Light travels through space as a wave, similar to an ocean wave.
Wavelength (l) (lambda) is the distance the light wave travels in one cycle.
Frequency (n) (Nu) is the number of wave cycles completed in each second.
Light travels at a constant speed (c) : 3.00×108 m/s.

26. Wavelength Versus Frequency
The longer the wavelength of light, the lower the frequency.
The shorter the wavelength of light, the higher the frequency.
Wavelength tells you the type of light. Frequency tells about wavelength-short or long. Amplitude tells you about the intensity of the light.

27. Exercise 1:
What is the frequency of light with wavelength of 500nm?
ln = c
n = c/l
n = 3.00 x 108 m/s / 500 x 10-9m
= 6.00 x 1014 cs-1 or Hertz
Exercise 2:
A photon (light) has a frequency of 6.0 x 104 Hz. What is the wavelength of this light in nm?
ln = c
l = c/n
l = 3.00 x 108 ms-1/6.0 x 104 cs-1
l = 5.0 x 103 m
l = 5.0 x 1012 nm

28. Light—A Continuous Spectrum
Light usually refers to radiant energy that is visible to the human eye.
The visible spectrum is the range of wavelengths between 400 and 700 nm.
Radiant energy that has a wavelength lower than 400 nm and greater than 700 nm cannot be seen by the human eye.

29. Radiant Energy Spectrum
The complete radiant energy spectrum is an uninterrupted band, or continuous spectrum.
The radiant energy spectrum includes many types of radiation, most of which are invisible to the human eye.

30. The Quantum Concept
The quantum concept states that energy is present in small, discrete bundles.
For example:
A tennis ball that rolls down a ramp loses potential energy continuously.
A tennis ball that rolls down a staircase loses potential energy in small bundles. The loss is quantized.

31. Bohr Model of the Atom
Niels Bohr speculated that electrons orbit about the nucleus in fixed energy levels.
Electrons are found only in specific energy levels (orbitals), and nowhere else.
The electron energy levels are quantized.

32. Emission Line Spectrum
When an electrical voltage is passed across a gas in a sealed tube, a series of narrow lines is seen.
These lines are the emission line spectrum. The emission line spectrum for hydrogen gas shows four lines.
Red, Green, Blue-Violet, Violet

33. Evidence for Energy Levels
Bohr realized that this was the evidence he needed to prove his theory.
The electric charge temporarily excites an electron to a higher orbit. When the electron drops back down, a photon is given off.
The red line is the least energetic and corresponds to an electron dropping from energy level 3 to energy level 2.

34. “Atomic Fingerprints”
The emission line spectrum of each element is unique.
We can use the line spectrum to identify elements using their atomic fingerprint.

35. Critical Thinking: Neon Lights
Most “neon” signs don’t actually contain neon gas.
True neon signs are red in color.
Each noble gas has its own emission spectrum, and signs made with each have a different color.

36. Energy Levels and Sublevels
It was later shown that electrons occupy energy sublevels within each level.
These sublevels are given the designations s, p, d, and f.
These designations are in reference to the sharp, principal, diffuse, and fine lines in emission spectra.
The number of sublevels in each level is the same as the number of the main energy level.
They classified the lines based on their visual appearance, and came up with the nomenclature s,p,d and f.Their meaning is
s – sharp
p – principle
d – diffuse
f – fine (or fundamental)
There are letters beyond this, which just follow the alphabet, so g,h,i etc.

37. Energy Levels and Sublevels, Continued
The first energy level has one sublevel designated 1s.
The second energy level has two sublevels designated 2s and 2p.
The third energy level has three sublevels designated 3s, 3p, and 3d.

38. Electron Occupancy in Sublevels
The maximum number of electrons in each of the energy sublevels depends on the sublevel:
The s sublevel holds a maximum of 2 electrons.
The p sublevel holds a maximum of 6 electrons.
The d sublevel holds a maximum of 10 electrons.
The f sublevel holds a maximum of 14 electrons.
The maximum electrons per level is obtained by adding the maximum number of electrons in each sublevel.

39. Electrons per Energy Level

40. Electron Configurations
Electrons are arranged about the nucleus in a regular manner. The first electrons fill the energy sublevel closest to the nucleus.
Electrons continue filling each sublevel until it is full, and then start filling the next closest sublevel.
A partial list of sublevels in order of increasing energy is as follows:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d …

41. Filling Diagram for Energy Sublevels
The order does not strictly follow 1, 2, 3, etc.
For now, use Figure 4.16 to predict the order of sublevel filling.

43. Writing Electron Configurations
The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel.
The sublevel is written followed by a superscript with the number of electrons in the sublevel. For example, if the 2p sublevel contains two electrons, it is written 2p2.
The electron sublevels are arranged according to increasing energy.

44. Writing Electron Configurations, Continued
First, determine how many electrons are in the atom. Bromine has 35 electrons.
Arrange the energy sublevels according to increasing energy:
1s 2s 2p 3s 3p 4s 3d …
Fill each sublevel with electrons until you have used all the electrons in the atom:
Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d 10 4p5
The sum of the superscripts equals the atomic number of bromine (35).

45. Quantum Mechanical Model of the Atom
An orbital is the region of space where there is a high probability of finding an electron.
In the quantum mechanical atom, orbitals are arranged according to their size and shape.
The higher the energy of an orbital, the larger its size.
All s orbitals have spherical shapes.

46. Shapes of p Orbitals
Recall that there are three different p sublevels.
All p orbitals have dumbbell shapes.
Each of the p orbitals has the same shape, but each is oriented along a different axis in space.

47. Location of Electrons in an Orbital
The orbitals are the region of space in which the electrons are most likely to be found.
An analogy for an electron in a p orbital is a fly trapped in two bottles held end to end.

48. Shapes of d Orbitals Recall that there are five different d sublevels.
Four of the d orbitals have a clover-leaf shape and one has a dumbbell and doughnut shape.

49. Orbital Diagrams Each box in the diagram represents one orbital.
Half-arrows represent the electrons.
The direction of the arrow represents the relative spin of the electron.

50. Orbital Diagrams
core notation for the electron configuration of the Na atom can be written as [Ne] 3S1
Energy levels start to overlap in energy (e.g., 4s is lower in energy than 3d.)

51. Electron Configurations
The way electrons are distributed in an atom is called its electron configuration.
Three rules:
1.Electrons fill orbitals first starting with lowest energy orbitals and moving upwards (Aufbau principle)
2. No more than two electrons can be placed in each orbital and no two electrons can fill one orbital with the same spin (Pauli exclusion principle)
3. For degenerate (equal energy) orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule).

52. Chapter Summary
Atoms are composed of protons, neutrons, and electrons.
The protons and neutrons are located in the nucleus, and the electrons are outside the nucleus.
Atoms are mostly empty space.
The number of protons is referred to as the atomic number for the atom.

53. Chapter Summary, Continued
All atoms of the same element have the same number of protons.
Isotopes are atoms with the same number of protons, but differing numbers of neutrons.
The mass number for an isotope is the total number of protons plus neutrons.
The atomic mass of an element is the weighted average of the masses of all the naturally occurring isotopes.

54. Chapter Summary, Continued
Light has properties of both waves and particles.
The particles of light are referred to as photons.
The energy of photons is quantized.
Electrons exist around the nucleus of atoms in discrete, quantized energy levels.
Electrons fill energy sublevels, starting with the lowest energy sublevel and filling each successive level of higher energy.