The Structure of the Atom- Chapter 4, 3

Unit 2 Atomic Structure & Nuclear Chemistry Section 1 Create your own notes! Obj.1 Section 2 Obj.2 Counting Subatomic Particles

The ______is defined as the smallest particle of an element that retains the properties of that element. Subatomic Particles & Their Properties Particle Symbol Location Relative Mass Actual mass in Electrical Charge in a.m.u* grams 1 amu 1.6710-24 g +1 p+

1 amu 1.6710-24 g 0 n0

.00055 amu 9.1110-28 g -1 e-

*a.m.u : atomic mass unit 1 amu (“atomic mass unit”) = 1.67  10-24 g

**The mass of the proton and neutron are very similar** **Their mass is about 2000 times larger than the mass of the electron**

What are the 2 regions of the atom as of now?

 Nucleus: small, dense center containing ______and ______.  Electron Cloud: region surrounding nucleus containing electrons and is mostly ______.

Counting Subatomic Particles Atomic Number Mass Number  Every atom has a different number of protons,  The number of protons and neutrons in a which is represented by the atomic number. nucleus of a specific atom is called the mass  The number of protons determines the identity number. of the atom Atomic Number = Number of protons 2 Options: Ex: Nitrogen’s atomic number is 7 [1] Add the number of given protons & neutrons together Mass # = protons + Calculating Electrons in a Neutral Atom (A.P.E) neutrons  A neutral atom has the ______number [2] USE ONLY when no element information of protons and electrons. has been given: Round atomic mass to a whole # protons = # electrons number to get an element’s most likely most common mass number. o Nitrogen’s atomic mass = 14.01 amu o Nitrogen’s mass number is  To calculate the number of neutrons, subtract ______the atomic number from the mass number  Number of NEUTRONS = mass number – atomic number Calculating Neutrons (M.A.N) Nitrogen: 14 – 7 = 7 neutrons  Overall Charge = protons - electrons Calculating Electrons in a Ion  An ion has a charge. The number of protons is Example: NOT EQUAL to the number of electrons. -1  Can be a positively charged ion called a [1] How many electrons does Br have? ______or a negatively charged ion called an ______

[2] How many electrons does Al+3 have?

Element Information: Nuclear Symbol LABEL THE PARTS BELOW

A A Z Z X X C C

SELF CHECK What are the number of protons, neutrons & electrons in each atom?

Hyphen Notation: LABEL THE PARTS!

Copper – 65 ______p __e __ n You Try: (Be Careful: Never use the given atomic mas on the periodic table to calculate mass number unless no other information has been provided. )

Nuclear Hyphen Atomic # Mass # Charge Proton Neutron Electron Symbol Notation Magnesium-25 +2

82 126 82 Section 2 Obj.2 cont. & Obj. 4 How Do Atoms Differ: Isotopes

 Isotopes are atoms of the same element that have the same number of protons but a different number of ______.  Most elements contain a mixture of 2 or more isotopes. Each one having its own mass and abundance.  Some isotopes are radioactive and unstable.  You can identify an isotope by its different mass number of the same element

Carbon-12 Carbon-13 12 13 C C 6 6

Fill in the chart using the diagram in the presentation Atomic Isotope Protons Neutrons Electrons Mass (a.m.u) number

Lithium-6

Lithium-7

Lithium-8

You Try ?

What the number of protons, neutrons and electrons in each? 19 18 203 194 F F Hg Hg

9 9 80 80 How to calculate Average Atomic Mass?

 Average atomic mass is the weighted ______of the masses of all naturally occurring isotopes. Average Atomic Mass Equation:

Average atomic mass = (% abundance of isotope x mass of 1st isotope) + (% abundance of isotope x mass of 2nd isotope) + ……… Cl-35 and Cl-37 Why is the AAM = 35.45 amu ?

Example 1: Element x has 2 natural isotopes. Calculate its average atomic mass. 1st isotope has a mass of 10.012 a.m.u. with 19.91% abundance. 80.09% of the 2nd element has a mass of 11.009 a.m.u.

Example 2: Calculate the average atomic mass of copper if it has 2 isotopes. 69.11% has a mass of 62.93 a.m.u and the rest has a mass of 64.93 a.m.u.

Section 3 Obj.3 Nuclear Decay Reactions

Nuclear chemistry is the study of the changes of the ______of atoms.  Nuclear Reactions involve changes within the nucleus where as chemical reactions involve the loss, gain or sharing of electrons.

The Nucleus

 Contains ______and neutrons. They are collectively called ______.

Radioactivity

 A ______nucleus holds together well. An unstable nucleus will decay or break down, releasing particles and/or energy in order to become stable.  An atom with an unstable nucleus is considered “______”.

Transmutation

 Type of nuclear reaction that will change the number of ______and thus will create a different ______. There are several ways radioactive atoms can decay into different atoms!

Basic Types of Radioactive Decay

Particle Symbol What Happens? Example Penetrating Power Type

Atomic number decreases by (1) LOW: Can be Alpha or ____ and mass number blocked by decreases by ____ paper/clothing

(100) MEDIUM: can penetrate the skin; Atomic number increases by need to be Beta or _____ but the mass number protected by *** stays the same clothing/thin metals like aluminum **A neutron becomes a ______and a high-speed electron that is discharged from the nucleus** (100000) HIGH: No change in atomic nor need to be Gamma mass number; occurs with protected by thick other types of decay concrete or metal like lead

Writing Balanced Nuclear Equations

The sum of ______and mass number must balance on both sides of the equation. Often problems will have 1 particle missing and you will need to identify it.

You Try! 1. Beta decay of zircomium-97 5. Complete this:

3. Alpha decay of americium-241 7.

8. 9. Making New Elements and Isotopes: Bombarding the Nucleus  All the transuranium elements (elements with atomic numbers higher than Uranium) have been made by bombarding the nucleus with neutrons and other atoms in accelerators

10. 11. 12. Section 4 Obj.3 cont. 13. Half Life

 Radioactive isotopes decay at a characteristic rate measured in ______.  A half-life time (T1/2) is the time required for ______of the amount of radioactive atoms to decay. The time ranges from ______to millions of ______. 14. HOW TO’s 15. 1. To calculate the number of half-lives, divide the half-life (T1/2) into the total time (T). Then cut the original amount in half the number of time determined by the # of half lives (HL). 16. 17. T / T1/2 = HL 18. 2. Algebraic equation to calculate the remaining amount left over after a certain number of half-lives have passed. 19. Amount remaining = (initial amount) (.5) HL 20. 21. Examples: 1. Suppose you have 20 grams of sodium-24. Its half-life is 15 hours. How much is left over after 60 hours. 22. 23. 24. 25. 2. Uranium-238 has a half-life of 4.46 x 109 years. How long will it take for 7/8th of the sample to decay? 26. 27. 28. 29. 30. You Try! 31. 1. 1.5 grams of a 12.0 g sample are left after 114 s. What is the half-life of radium-222? 32. 33. 34. 35. 36. 2. A sample of 3x107 Radon atoms is trapped in a basement that is sealed. The half-life of Radon is 3.83 37. days. How many radon atoms are left after 31 days? 38. 39. 40. 41. 42. 43. Section 5 Extension 44. Fusion & Fission 45. 46. Nuclear Fission  Large atoms ______into smaller atoms

 Generates huge amounts of ______. 47.  Carried out in ______ Could result in a chain reaction of fission like the

______48. 49. 50. 51. 52. 53. 54. Nuclear Fusion  Smaller atoms are ______to form a larger atom.  Occurs in the ______and ______ Generates huge amounts of ______55.