Atomic Structure and Nuclear Chemistry

1. Atomic Structure and Nuclear Chemistry Chapter 4 and 18

2. Elements a.k.a atoms • Robert Boyle first defined an element as a substance which could no longer be broken down into other substances • Each element has unique properties • Many early scientists speculated how the element or atom was structured • Theory of atomic structure evolved from early thoughts to today’s atom

3. What do you know about the atom? • Take a moment and create a concept web about the atom. • Work on your own. • You have about 5 minutes. • Jot down everything you can connect to atoms. Atoms

4. Timeline of the Atom Democritus (460-370 B.C.) understood that if you cut a stone in two pieces, each piece contained the same material as the original stone. He also believed that you could do this an infinite number of times. He called these infinitesimally small pieces of matter atomos, meaning "indivisible.“ Rutherford’s gold foil experiment leads to atomic nucleus and in 1919 the introduction of the proton 1808 - Dalton first proposed a theory on atoms Discovery of the electron by J.J. Thompson in the late 1890’s 1910 -Lord Kelvin’s “plum pudding theory 1912 –Bohr model of the hydrogen atom Mid-1920’s – wave mechanical model 1932 – Rutherford and his coworker Chadwick identified the neutron

5. Who started… John Dalton (1766 – 1844) was an English scientist who made his living as a teacher in Manchester.

6. Dalton’s Atomic Theory (p.88) • Elements are composed of atoms • All atoms of a given element are identical • Atoms of a given element are different from those of any other element • Atoms of one element combine with atoms of other elements to form compounds • Law of Constant Composition: all samples of a compound have the same proportion of the elements as in any other sample of that compound • Atoms are indivisible in a chemical process. • all atoms present at the beginning of a chemical process must also be present at the end of the process. • atoms are not created or destroyed, they must be conserved. • atoms of one element cannot be turned into atoms of another element

7. Atomic Structure History • Discovery of the Electron • 1st atomic particle identified • In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. • Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. This creates a beam of negatively charged particles bent by an electric field.

8. Conclusions from the Study of the Electron • Cathode rays have identical properties regardless of the elemental gas used to produce them. • Therefore, all elements must contain identically charged particles (electrons). • Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons • Electrons have so little mass that atoms must contain other particles that account for most of the mass • An electron is a tiny, negatively charged particle • The next step…. What is the positive charge?

9. Models of the atom

10. William Thomson’s (Lord Kelvin’s) Atomic Model Lord Kelvin believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model (easier to think of as “chocolate chips" in chocolate chip cookie dough.)

11. Rutherford’s Gold Foil Experiment

12. Rutherford’s Gold Foil Experiment • Rutherford shot α (alpha) particles at a thin sheet of gold foil (think: bullet = alpha particles, target atoms = gold foil) • α particles are positively charged • gold atoms are about 50 larger than α particles. • Particles were fired at a thin sheet of gold foil • Particles hit on the detecting screen (film) were recorded

13. (a) The results that the metal foil experiment would have yielded if the plum pudding model had been correct. (b) Actual results known as Rutherford’s model. Over 98% of the  particles went straight through About 2% of the  particles went through but were deflected by large angles About 0.01% of the  particles bounced off the gold foil Most of the volume of the atom is empty space

14. Rutherford’s Conclusion: A Nuclear Model The atom contains a tiny dense center called the nucleus • the volume is about 1/10 trillionth the volume of the atom The nucleus is essentially the entire mass of the atom (extremely dense) The nucleus is positively charged • the amount of positive charge of the nucleus balances the negative charge of the electrons The electrons move around in the empty space of the atom surrounding the nucleus

15. Finally, the neutron.. • Discovered in 1932 by Chadwick based on the idea from Rutherford • Has no charge • Is located in the nucleus • Mass a mass of 1 amu (actually, it’s slightly larger than a proton but for our work the mass is the same)

16. So, the structure of the atom ……. • The nucleus was found to be composed of two kinds of particles • Some of these particles are called protons • charge = +1 • mass is about the same as a hydrogen atom • Since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons • The other particle is called a neutron • has no charge • has a mass slightly more than a proton

17. The Modern Atom • We know atoms are composed of three main atomic particles - protons, neutrons and electrons • The nucleus contains protons and neutrons • The radius of the atom is about 100,000 times larger than the radius of the nucleus A nuclear atom viewed in cross section.

18. Summary of Atomic Particles Going beyond the electron, proton, and neutron

19. Describing an Atom How many protons, neutron, and electrons does an atom have?

20. Atomic Structure - protons The number of protons in an atom of a given element is the same as its atomic number (Z). • (Z) is found on the Periodic Table, whole # for each element

21. Atomic Structure - neutrons • Mass number = protons + neutrons; always a whole number. • # of Neutrons = mass number - # of protons • Atomic mass – decimal number in each element’s box on the periodic table. If you round the atomic mass of an element to the closest whole number you generally get the mass # for that element.

22. Atomic Structure - Electrons • # of Electrons = # of protons if the atom is neutral • If the chemical symbol is written with a charge, representing an ion, the charge indicates the number of electrons that have been added or removed from the atom. • If the ion has a positive charge (cation), subtract that charge from the # of protons to get the number of electrons. • If the ion has a negative charge (anion), add that charge number to # of protons to get the number of electrons. • # of Electrons = # protons – charge • Charge = # protons - # electrons

23. Representing atomic particles in atoms • The number of each type of atomic particle (proton, neutron, electron) is determined by using symbols. • There are several different ways to write an element: • Atomic symbols • Nuclear symbols • Atomic Symbols include the element symbol and a charge if any. • C – neutral carbon • C+4 – carbon cation • C-4 – carbon anion

24. NuclearSymbols Mass number (p+ + no) Charge (if any) Cl-1 –38 Fluorine-18 U 238 + Element symbol 92 Atomic number (number of p+) Element name Mass number Element symbol with charge Mass number

26. Isotopes atoms of an element with the same number of protons but different numbers of neutrons

27. Two isotopes of sodium.

28. Isotopes Examples • 3517Cl 3717Cl • H-1 H-2 H-3 Copper – 63 Copper – 65

29. Isotopes and Atomic Mass

30. AMU • When we think about the mass of an atom, we use atomic mass units (amu). • A proton is 1 amu • A neutron is 1 amu • Add up protons and neutrons to get the mass number (not atomic mass) Why? • Most elements in nature have isotopes • All these isotopes contribute to the average atomic mass (listed on the table)

31. Determining Average Atomic Mass • The average atomic mass seen on the periodic table is a combination of all an element’s isotopes and their abundance. • To determine the average atomic mass for an element, you must • Multiply the percentage (percent abundance) of each isotope of the element by its mass number. • Add the products of the multiplications together. • Divide by 100. • Your answer should be very close to the atomic mass of the element for that element

32. Average Atomic Mass Examples • Find the average atomic mass of each of the following elements from their percentages and mass numbers. • 69.17% 63Cu and 30.83% 65Cu • 5.85% Fe-54, 91.75% Fe-56, 2.12% Fe-57 and 0.28% Fe-58

33. Nuclear Reactions • What you just did was write nuclear reactions. • Typical reactions are decay reactions and capture reactions. • What did you notice about the products of the reactions? • The products of the reactions are isotopes of the element. • Nuclear reactions produce different particles that are not elements. • You need these particles to balance out the protons and neutrons in the nuclei. • In a nuclear reaction, the atomic number (Z) and the mass number (A) are conserved.

34. Radioactive decay • Radioactive decay is a natural process. • Only a handful (~200) of the known isotope nuclei (2000) do not decay • A nuclei may spontaneously kick out a particle, forming a new element. • There are four common particles: • The alpha particle - α (gold foil) • The beta particle – β • The gamma particle – γ • The positron particle • In many cases, the process does not stop at one step, but rather a combination of steps. This is known as a decay series.

35. Capture Reactions

36. Nuclear Transformations • Nuclear transformation is the changing of one element to another (modern alchemy!!!) using larger atoms • Rutherford observed 1st transformation in 1919 • Marie Curie and her husband another transformation (14 years later) • Electron capture is the “natural” nuclear transformation • Man-made elements are made by bombarding two nuclei together

37. The nuclear particles

38. Radioactivity Review • Radioactivity is expressed as either a decay process or a capture process • Decay processes can be connected as a chain • Nuclear transformations use larger atoms to create different elements than what you started with

39. Review of nuclear symbols • Mass number is in upper left of symbol • Atomic mass is in lower left of symbol • Nuclear particles are written in the nuclear format U 238 92

40. The alpha particle (α) • The alpha particle is actually a helium nucleus. • It is the weakest of the decay particles. • A alpha particle has a mass number of 4 and an atomic number of 2. • If an α particle is added to a nuclei, the mass number will increase by 4 and the atomic number by 2. • If an α particle is released from a nuclei, the mass number will decrease by 4 and the atomic number by 2. • For example:

41. The Beta particle (β) • The beta particle is an electron. • It has no mass. • It does have a charge • Examples include: • Sometimes a nucleus will grab an electron that is close. This is called electron capture.

42. The Gamma Particle (γ) • This is also know as the gamma ray. • This is a high energy photon of light. • Picked up by specific detectors. • It is the strongest (most dangerous) of the decay particles. • Example:

43. The Positron • The positron is similar to a beta particle, but has a positive charge. • Example:

44. Neutron Emission • Neutron emission or decay does not change the element, only the mass • Example:

45. Balancing Nuclear Equations • Mass # and the atomic # totals must be the same for reactants and the products. • 3919K  3517Cl + ___ • 20682Pb  0-1e + ___

46. Writing Balanced Nuclear Equations • Alpha decay of Cu-68 • Gamma emission of Thorium-235 • Positron emission of P-18 • Astatine-210 releasing 3 neutrons • Electron capture with Ti-45

47. Nuclear Chemistry and half-life Balancing Nuclear Equations Half-life

48. Half-Life and Nuclear Stability • Radioactive isotopes or nuclides all decay because they are unstable, some just breakdown much faster than others • Geiger counter or scintillation counters are used to detect particles. • Half-life – amount of time for half of the original sample to decay • For two samples of the same isotope, regardless of the sample size, after one half-life, only half of the original amount of sample remains.

49. Sample Half-lives • IsotopesHalf-Live • Carbon – 14 5730 years • Sodium – 24 15 hours • Bismuth – 212 60.5 seconds • Polonium – 215 0.0018 seconds • Thorium – 230 75400 years • Thorium – 234 24.1 days • Uranium – 235 7.0 x 108 years • Uranium – 238 4.46 x 109 years

50. Working with half-life • A material has t1/2 = 10 minutes. If you begin with 16g, how long will it take to decay to 2 g? • Begin with 16 g, 1 half life gets you to 8 g. • 2 half lives get you to 4 g. • 3 half lives get you to 2 g. • So, 3 x 10 minutes = 30 minutes. • A material has t1/2 = 150 years. If you begin with 100 g, how long will it take to decay to 3.125 g?