2. Unit 3 Chapter 3 & Chapter 11 The Structure of the Atom & Nuclear Chemistry

3. Early Theories of Matter Democritus (460-370 B.C.) proposed & believed that 1.Matter was not infinitely divisible 2.Made up of tiny particles called atomos 3.Atoms could not be created, destroyed, or further divided - This sounds like the beginning of…..

4. John Dalton 19 th century Dalton revised Democritus's ideas based upon the results of scientific research he conducted Dalton’s atomic theory

5. Dalton’s Atomic Theory Elements are made of extremely small particles called atoms All atoms of a given element are identical  Same size, mass and chemical properties Atoms are indivisible in chemical processes.  They can neither be created nor destroyed in a chemical reaction  A chemical reaction simply changes the way the atom is grouped together A compound has a constant composition of its elements  Different atoms combine in simple whole number ratios to form compounds

6. Dalton’s Atomic Theory Dalton’s theory was supported by 2 Laws: The Law of Conservation of Matter (Lavoisier) The Law of Definite or Constant Composition (Proust) Not all of his theory is correct!!

7. What is an atom? An atom is the smallest particle of an element that retains the properties of the element

8. The Nuclear Atom Ernest Rutherford developed a model of the atom. His model consisted of the following ideas:  An atom consists mostly of empty space through which electrons move  Electrons are held within the atom by their attraction to the positively charged nucleus  Tiny, dense region called the nucleus, which is located in the center of the atom Nucleus contains all of an atom’s positive charge and virtually all of its mass

10. Three Subatomic Particles Electron Proton Neutron

11. Subatomic Particles in the Nucleus Positively charged particles called protons Proton’s charge: equal to but opposite that of an electron Why don’t the protons repel each other in the nucleus? Neutral particles called neutrons Neutron mass nearly equal to that of proton Neutron has no charge

12. Outside the nucleus Negatively charged particles are called electrons  move through empty space in atom  Negatively charged

13. Atomic Number The number of Protons in the nucleus Examples:  Carbon (C) has 6 protons Atomic number is 6  Copper has 29 protons Atomic number is 29

14. Atomic Number In uncharged atom, atomic number is also the number of electrons  Why?  If an atom is charged, then it is an ion Uncharged atom: Atomic number = # of protons = # of electrons Take out your periodic table

15. Mass Number To find the Mass number # protons + # neutrons To find # neutrons mass number – proton (or atomic number) Mass numbers are always WHOLE #’s!!

16. Symbols for Atoms X= symbol of element A= mass number Z= number of protons X or X AZAZ A

17. Isotopes and Mass Number C C carbon-12 carbon-13 Isotopes are atoms with the same number of protons but different number of neutrons 13 6 12 6

18. Isotopes and Mass Number Example: 3 types of Potassium All 3 types contain 19 protons and __ electrons # of Protons # of NeutronsMass Number 1920 1921 1922

19. What’s the difference between mass number and average atomic mass? C carbon-12 But if you look on the periodic table, the number states 12.01… 12 6

20. Mass of Individual Atoms Mass of protons and neutrons: 1.67 x 10 -24 g Mass of electron is 1840 times smaller than that of protons and neutrons 9.11 x 10 -28 g

21. Atomic Mass Unit (amu) Small mass #’s are not easy to work with, so the atomic mass unit (amu) was developed One atomic mass unit (amu) is equal to 1/12 the mass of a carbon-12 atom The mass of 1 amu is nearly equal to the mass of one proton or neutron

22. Mass of Individual Atoms Atomic mass:  The weighted average mass of the isotopes of an element  Example: Chlorine Mixture of 75% chlorine-35 and 25% chlorine-37 Atomic mass = (0.75)*35 + (0.25)*37 = 35.5 amu

23. What’s the difference between mass number and average atomic mass? Mass number- specifically about one isotope Average atomic mass- includes the masses of all the different isotopes for that atom

24. Nuclear Reactions Chemical reactions  What can NOT change in a chemical reaction? Nuclear Reactions: changes that occur in the nucleus of an atom  The nucleus is unstable!! Most atoms have unstable nuclei.

25. Why are some stable, while others are not? Primary Reason: ratio of the neutrons to the protons (n/p) An atom is most stable when the ratio is 1:1 The maximum ratio of stability is around 1.5 : 1

26. Radioactivity An unstable nucleus emits rays and particles, called radiation, to become stable  The process is called radioactivity  Gain stability by LOSING energy

27. Radiation Three types of radiation produced:  Alpha  Beta  Gamma

28. 3 Radioactive Particles 1. Alpha Particle (+) oIt travels about 1/10 the speed of light (slowest) oIt is the largest, most massive particle oIt is the most dangerous if ingested oIt has the least penetrating ability - paper can stop this particle

29. 3 Radioactive Particles 2. Beta Particle (-) oFast accelerated electron oEjected when a neutron is converted to a proton in the nucleus oTravels 1/4 the speed of light. oIt is lighter and faster than the alpha particle. oAverage penetrating ability - can be stopped by heavy clothing

30. 3 Radioactive Particles 3. Gamma Ray oNot really a particle; it is a form of energy from the electromagnetic spectrum (EM) oHas no mass and no charge. oAlways accompanies either beta or alpha radiation. oVery high powered oTravels at the speed of light. oHighest penetrating ability - can be stopped by heavy shielding such as lead.

31. Alpha Decay An alpha particle is released A new atom is formed in the product in which: a) the atomic number is lowered by 2 b)the mass number is lowered by 4 Example:

32. Beta Decay A neutron is converted in the nucleus of an atom resulting in: a) beta particle released b) a new atom in the product whose atomic number increases by 1 c) mass number does not change Example:

33. Gamma Radiation It is the “Energy” of the reaction Always accompanies alpha and beta emissions. Regardless of radiation or decay, The Law of Conservation of Matter must be observed!! The nuclear equation must be BALANCED!!

34. RADIOACTIVE PARTICLES Summary

35. Half Life The half life of a radioisotope is the time it takes for half of the sample to decay. Half-life is represented by “t 1/2 ” The number of half-lives that have passed is represented by “n”

36. Amount Remaining You can use two equations, depending on what information you have If you have the # of half-lives that have passed.. initial amount x (1/2) n n= # of half-lives that have passed If you DON”T have the # of half- lives that have passed.. initial amount x (1/2) P/ t1/2 p= time that has passed t 1/2 = time of half-live

37. Half-life problem If the half life element A is 3 hours and you have 90 grams of it, how many grams would be left after 9 hours? What percent of it is remaining?

38. Handout #5 Equations to Add To find the half-life  t 1/2 = time passed/ # of half-lives passed To find the # of half lives that have passed…  n= time passed/ one half-life Total time passed= n times t 1/2

39. Problem 3 Given information  I= 64g  r= 2.0g  p=12.5 hours  n=?  One half-life=??

40. FUSION and FISSION Fission is the splitting of nuclei  resulting in a tremendous release of energy  heavy atoms split so they can become more stable Fusion is the combining of nuclei  resulting in more energy being released than fission  The sun produces energy as a result of nuclear fusion  2 H atoms combine to form Helium