1 Chapter 4: Objectives: the student will be able to: compare and contrast early atomic models; distinguish between subatomic particles in terms of mass and charge; dissect the atom into its component parts; explain the role of atomic number in determining the identity of an atom; define an isotope and explain why atomic masses are not whole numbers; calculate the number of protons, electrons and neutrons in an atom given the mass number and atomic number; explain the relationship between unstable nuclei and radioactive decay; characterize alpha, beta, and gamma radiation in terms of mass and charge; interconnect the changes in the theory of the atom with the idea of the scientific method

2 Dalton’s Atomic Theory (1808 A.D.)-very similar to ideas of Democritus (~400 B.C.) Each element is made up of tiny particles called atoms that can not be broken down into smaller particles. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. Chemical reactions involve reorganization of the atoms - changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

3 Cathode Ray: stream of charged particles produced by an electric field. The green “glow” in the picture. The particles move from the cathode (negatively charged plate) to the anode (positively charged plate).

4 Cathode “rays” are influenced by both magnetic and electric fields (electric field shown in the diagram). Opposite charges are known to attract each other. Since the cathode “ray” was deflected towards the positively charged plate, what must the charge be of the particles in the cathode “ray”?

5 The negatively charged particles produced by the electric field were named electrons. J.J Thompson (~1890)-measured the charge to mass ratio of the electron. He discovered that an electron must be smaller than the smallest known atom. Whoops! There goes part of Dalton’s theory of the atom-an atom is not the smallest particle of matter. However, Dalton was correct that an atom is the smallest particle of matter that can still be recognized as a specific element.

6 Robert Millikan (1909)-measured the charge of an electron. Knowing the charge to mass ratio of an electron from Thompson, Millikan can use the charge he has measured to calculate the mass of an electron! Mass of e  1 = 9.11X10  28 gMass of H atom = 1.676X10  24 g (1.676X10  24 g)/(9.11X10  28 g) = 1840 A Hydrogen atom is 1840 times more massive than an electron.

7 After Thompson’s and Millikan’s work, a new “model” of an atom must be produced. In the old one an atom was much like a ball bearing because it was the same throughout. The new model must: 1) allow electrons to be removed from the atom; 2) must explain why an atom is neutral; 3) must explain where most of the mass of an atom is located. Plum Pudding Model: Electrons are evenly distributed throughout a uniform positive charge equal to the negative charge of all of the electrons. The part with the positive charge is where most of the mass of an atom is located.

8 Ernest Rutherford (1911) “gold foil” experiment Large, high energy particles (alpha particles, +2 charge) do not always pass straight through a thin sheet of matter. The alpha particle being deflected or reflected is like having a bullet bounce off of a sheet of tissue paper.

9 Rutherford’s experiment does not fit with the Plum Pudding Model (a) because not all of the alpha particles do not pass straight through. Rutherford proposed a nuclear model of the atom. The small, dense nucleus contains virtually all the mass of the atom and all of the positive charge while the negatively charged electrons exist apart from the nucleus (b). How does this model fit Rutherford’s results?

10 Rutherford called the positive particle in the nucleus a proton. In 1932, James Chadwick discovered a neutral particle in the nucleus and called it a neutron. Summary of Subatomic Particles (mid 1900’s) Relative Mass 1/1840 ~1

11 If all atoms contain the same types of particles, what makes one atom hydrogen and another carbon? Nuclear Model of an Atom The number of protons in the nucleus determine what element an atom is. Remember that in normal atoms, the number of protons is equal to the number of electrons (they are neutral-a total charge of zero). How many times bigger is the atom compared to the nucleus? ______________ ________________________________________________________________ Small, very dense nucleus containing massive protons and neutrons, surrounded by small rapidly moving electrons

12 Atomic Number: Atomic number is the number of protons in the nucleus of an element. Since in neutral atoms the number of protons is equal to the number of electrons, atomic number also indicates the number of electrons in an element. The Periodic Table is arranged according to atomic number. For example, hydrogen is element 1 and has 1 proton in its nucleus. Similarly, helium is element 2 and has 2 protons in its nucleus. How many protons does element number 25 contain? What is the name of this element? How many electrons does element 25 contain? What is the charge of one atom of element 25?

13 Isotopes are atoms with the same number of protons (same element) but with a different number of neutrons. Na-23 ( = 23) Na-24 ( = 24)

14 Mass Number: Mass Number is the number of protons plus the number of neutrons in the nucleus of an atom. Mass # = # Protons + # Neutrons Remember: Atomic # = # Protons Therefore: Mass # = Atomic # + # Neutrons If you have Mass # and Atomic #, you can find # Neutrons. From Algebra: a = b + c therefore if you are given a and b, you can find c using this equation: c = a  b

15 Element Symbol Mass Number Atomic Number # Neutrons = Mass #  Atomic # or # Neutrons = A  Z Since: Mass # = Atomic # + # Neutrons Remember that the mass number must be the largest number in the problem! X-A is another way to identify an isotope Identifying Specific Isotopes

16 Examples: Three Isotopes of Carbon (atomic number = 6) How many Protons, Electrons, and Neutrons are in each atom? Protons Electrons Neutrons 666From Atomic # Mass #  Atomic # C-12C-13C-14

17 Use a periodic table to help you complete the following table.

18 AMU: Atomic Mass Unit is a unit of mass based on 12 C. By Definition: the mass of one atom of 12 C = 12 amu, so 1 amu = 1/12 of the mass of one atom of 12 C On this scale, 1 proton = amu 1 neutron = amu 1 electron = amu The Mass of “C” on the periodic table is amu because pure carbon is made up of both 12 C (12 amu exactly) and 13 C (close to 13 amu). The mass of “C” is a weighted average of all isotopes of C!

19 Weighted Average: Average Mass = [(% abundance of isotope one)*(mass of isotope one) + (% abundance of isotope two)*(mass of isotope two)] Average Mass of Mg: (78.99%)*( amu) + (10.00%)*( amu) + (11.01%)*( amu) amu 24 Mg 25 Mg 26 Mg IsotopeAbundanceMass 78.99% 10.00% 11.01% amu amu amu Data for the known isotopes of Magnesium amu

20 Estimating the abundance of isotopes in an element. The atomic mass (numbers in black beneath the element symbols on the periodic table in the classroom) are weighted averages of all stable isotopes of that element. If we know the atomic mass we can make a guess about which isotope may be most abundant. Example: Carbon is amu. Since this is very nearly 12, we would guest that most of the carbon atoms in the world are C-12 and that there is very little of other isotopes. This method will not always work, but it can be used to get an idea. If we know which isotopes are present in an element and its atomic mass, we can start to estimate the abundance of the isotopes. Example: Cl-35 and Cl-37 are the only stable isotopes of chlorine. The atomic mass of Cl is amu. Since is closer to 35 than to 37, there is going to be a greater abundance of Cl-35 than of Cl-37. Using math to estimate: 50%* %*37 = 36 and 75%* %*37 = 35.5

21 + Beta Decay (  -particle is produced) A high energy electron is a  -particle ++ Alpha Decay (  -particle is produced) A helium nucleus is an  -particle High energy “light” is  -radiation Gamma Radiation (  -ray is produced) Types of Radioactive decay: occur because the nucleus of an isotope is not stable.

22 Alpha Radiation has +2 charge and mass of 4 amu Beta radiation has  1 charge and mass of (1/1840 amu) Gamma radiation has 0 charge and mass of 0 amu (+) ()() Electrically charged plates Lead Block Source of radiation is inside. target screen Why should we expect that beta would bend more than alpha?

23 Radioactive decay is dangerous because the particles have very high energy compared to “normal” matter, and it damages “normal” matter when it hits it. Gamma has the most energy and is the hardest to block out. It can penetrate several feet of concrete because it has no mass (it is really just a high energy form of “light”). Beta has less energy than gamma but more that alpha. Because of its small mass, it can penetrate more matter than alpha does. Alpha has the least energy of the three and can be blocked by just an inch or two of concrete. It has a large mass and will “hit” other atoms more easily than the much smaller beta particle. The danger in being exposed to radiation increases as you are exposed to more and more of it. Too much can cause death quickly. Smaller amounts may cause cancer or other diseases.

24 Since nuclear reactions create and destroy atoms, we do not balance the number of atoms on both sides of the equation. Nuclear Equations must balance mass number and atomic number! + + Examples: Radioactive decay is a nuclear reaction and is not like chemical reaction!

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