RSPT 1060

1. RSPT 1060 MODULE B: Basic Chemistry Lesson #1 Atomic & Subatomic Matter

2. Why Chemistry? • Respiratory Therapists must have a basic knowledge of the principles of chemistry … • To better understand the functioning of the human body • To better appreciate the clinical concepts of: • Arterial blood gas interpretation • Fluid and electrolyte physiology • Nutrition • Pharmacology

3. Objectives • At the end of this module, the student will: • Define terms associated with atomic and sub-atomic matter. • Differentiate between the types of matter. • Describe what each item in an element’s box on the periodic table represents. • Compare the composition of the elements of the universe, the earth’s crust and the human body. • Differentiate between an atom, elements, molecules and compounds.

4. Objectives • At the end of this module, the student will: • Describe the components of an atom and the purpose of each. • Differentiate between atomic number, atomic mass and mass number. • Explain what an isotope is. • Explain what determines physical and chemical properties of an element.

5. Matter • What is Matter? • Anything that • Takes up space • Has mass (weight), and • Can be perceived by the senses. • If it’s “something” it’s matter, if it’s “nothing”, it’s not matter • The primary states of matter are: • Solid • Liquid • Gas

6. Divisions of Matter Matter Mixture (Homogeneous or Heterogeneous) Pure Substance(Homogeneous) Elements Compounds Solution Suspension Colloid

7. Matter - Pure Substances • Matter in it’s simplest form. • Atom, Element, Molecule, Compound • Always the same regardless of where it is found. • Oxygen (O), water (H2O), table salt (NaCl) • It cannot be broken down any further without a chemical or nuclear reaction. • It will then become a different substance. • Uranium in a nuclear bomb • Pure substances are homogenous. • Uniform in structure or composition throughout .

8. Matter • So…what is an element or compound?

9. Elements • An element is a pure form of matter. • Other pure forms of matter include: • Atoms • Molecules • Compounds

10. Elements • Large collection of atoms of the same type. • Substance that cannot be broken down further and still maintain its identity. • All atoms have same atomic number. • Not bonded together, only existing together. • A listing of all the elements known to man is called the Periodic Table.

11. Introduction to the Periodic Table 117 Elements 1 Hydrogen H 1.01 Atomic Number Element Name Symbol or Abbrev. Atomic Mass Unit (AMU) http://www.ceet.niu.edu/mrdl/software/Periodic%20Table.htm

12. Elements of the Universe 91% of all atoms are Hydrogen 9% of all atoms are Helium The other 115 elements are found in traces. 

13. Elements of the Earth’s Crust 60.1% = oxygen 21.1% = silicon 6.1% = aluminum 2.9% = hydrogen 2.6% = calcium 2.4% = magnesium 2.2% = iron 2.1% = sodium

14. Elements of the Human Body

15. Atoms • Smallest “particle” of an element.

16. Molecule • Smallest “particle” of a pure substances (element or compound) bonded together. • Combination of similar atoms (O2 - element) • Combination of different atoms (H2O - compound)

17. Compound • Substance composed of a large collection of molecules. Can be broken down by chemical means into molecules or elements. • Often will have properties unlike those of its constituent elements.

18. Pure Substances A B

19. The Atom • Smallest particle of an element which still maintains the chemical properties of the element. • Head of a pin could hold 100 trillion atoms.

20. The Atom • If broken down further by a nuclear reaction, an atom would become particles: • Electrons • Protons • Neutrons.

21. The Atom • If broken down further, protons and neutrons are made of subatomic particles: • Positrons • Mesons • Neutrinos

22. The Helium Atom Nucleus Proton 2 Protons (+) and 2 Neutrons (No Charge) 2E- Neutron Smallest particle of an element. Electron

23. Atom - Composition • Nucleus • Proton (+) nucleon • Neutron (No charge) nucleon • Electron cloud or shell • Electron (-)

24. Atom - Nucleus • The nucleus is the small, dense positively charged center of the atom • It contains protons and neutrons (nucleons) • The nucleus only comprises 1/100,000 of the size of the atom even though it is constitutes the vast majority of the atom’s mass.

25. Atoms - Nucleus • Nucleons • Protons • One Proton is 1836 times the size of an electron • The number of protons determines the atomic number. • The number of protons is always equal to the number of electrons • This allows for a neutral charge of the atom. • Neutrons • The number of neutrons can vary • The number of neutrons determines the number ofisotopes an element will have. • Isotope: One of two or more atoms having the same atomic number but different mass numbers.

26. Atom - Electrons • 99.9% of the atom is open space where the electrons travel (electron cloud or shell) • 99.99% of an atom is the negatively charged electron cloud • This cloud actually determines the size of the atom

27. Atom - Electrons • Electrons do not contribute to the mass of the atom; only the size. EXAMPLE: If the electron cloud was the size of Ford Field, the nucleus would be smaller than a pea at the center of the field. • The nucleus determines the mass.

28. Atom – Size & mass • Electrons determine size • Nucleus (protons & neutrons) determine mass

29. Atom – Electron Number and Arrangement • The number and arrangement of the electrons determine the chemical properties of an element. • How it acts in relation to other elements • How it acts in a chemical reaction

30. The Periodic Table 112 Elements 1 Hydrogen H 1.01 Atomic Number Element Name Symbol or Abbrev. Atomic Mass Unit (AMU)

31. Atomic Number • The number of protons in the atom of a given element. • All atoms of an element have the same number of protons and electrons. • This never changes. • Because atoms are neutral, the atomic number also indicates the number of electrons. EXAMPLE: Boron has an Atomic # 5 • This means there are _____ protons & _____ electrons 5 5

32. Atomic Mass Unit • Abbreviated as (amu). • Reflects the mass of the most frequently found form of an element in nature. • The unit amu is a unit of measure made up by scientists. • It is used as a unit of measure for a particle that is extremely small. • 1 amu = 1.6606 x 10-24 grams

33. Mass of an Atom and the amu • The mass of an atom is too small to express in grams • Hydrogen atom = 1.7 x 10-24 gram. • The relative scale of atomic mass units is used instead of grams & scientific notation.

34. Comparative Example • 12 eggs = One dozen • Dozen is a unit of measure made up by farmers. (not really) • Dozen is a simple unit of measure that represents a larger number (12)

35. Mass of an Atom • Mass is composed mainly of the mass of protons & neutrons • Proton = 1 amu • Neutron = 1 amu

36. Carbon and the amu • All elements are compared to the mass of carbon. • 1 amu = 1/12 the mass of a Carbon atom • Carbon has 6 protons & 6 neutrons • It’s atomic mass is 12.011 AMU • Carbon is a point of reference for all other elements • Hydrogen is 1/12 the mass of carbon so it has a mass of 1 amu • 1 proton & 0 neutron • Magnesium has twice the mass of carbon so it has a mass of 24 amu • 12 protons & 12 neutrons

37. Isotope • There may be different forms of atoms of the same element. • This occurs when the number of neutrons varies. • Atoms of the same element with differing numbers of neutrons are calledisotopes

38. Isotopes and Physical Properties • Neutrons will determine the physical properties which vary slightly between isotopes. • Result:The same element may “appear” slightly different depending on which isotope you look at. All isotopes should “act” the same because the electron numbers don’t change. • Only 20 elements exist without isotopes.

39. Isotopes and Medicine • We hear about isotopes most often in nuclear medicine. • Body scans use isotopes (Xenon) • Ventilation & perfusion of the lungs • Bone scans • Radioactive material is injected in the blood or inhaled into the lungs • Image forms on radiology film showing areas that isotope has been exposed to

40. Key Facts about Isotopes • Isotopes: • Atoms of the same element • BUT, Have different numbers of neutrons • Atomic number on periodic table does not change (Same # of Protons) • Atomic Mass (amu) on periodic table does not change (Average of most common isotopes) • Mass numberchanges (Actual number of protons and neutrons)

41. Example of an Isotope: Chlorine Example: Chlorine #17 Atomic mass 35.45 • Most common form (76% of the time) • Cl-35 with a mass of 34.97 amu • Less common form (24% of the time) • Cl-37 with a mass of 36.97 amu Calculation: • (0.76)(34.97) = 26.5772 amu • (0.24)(36.97) = 8.8727 amu 35.45 amu (average on periodic table)

42. Isotope Chlorine: • Atomic mass – Atomic # = the average #neutrons • 35 – 17 = 18neutrons in most common form

43. Isotopes and Mass Number • “Mass Number” • Each isotope has its own mass number. • Not on the periodic table • Is the actual total number of protons + neutrons • The number of neutrons can change so the mass number can change. • Protons + Neutrons = Mass number

44. Isotope Example: Chlorine #17 Atomic mass 35.45 • Most common form 76% of the time • Cl-35 with a mass of 34.97 amu • Mass # 35 • Mass # - Atomic # = Neutrons • 35 – 17 = 18 • Less common form 24% of the time • Cl-37 with a mass of 36.97 amu • Mass # 37 • Mass # - Atomic # = Neutrons • 37 – 17 = 20

45. Example: Potassium (K) • Atomic Number = 19 • Mass Number = 39 • Mass Number = 40 • Protons = _______ • Electrons = _________ • Neutrons = Mass # – Atomic # = ___________ or __________

46. Oxygen Isotopes (AMU 15.9994)

47. ASSIGNMENTS • Read: Chemistry Book to assist in completing the objectives. • Self-Assessment