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 Pure Substance (Homogeneous) Elements Compounds Solution Suspension Colloid Mixture (Homogeneous or Heterogeneous)

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 (H 2 O), 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 Atomic Number Element Name Symbol or Abbrev. Atomic Mass Unit (AMU) 1 Hydrogen H 1.01 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 Major Elements WaterMacronutrientsMicronutrients Hydrogen Oxygen Carbon Nitrogen 66% of the body. Calcium Chlorine Potassium Phosphorus Magnesium Sulfur Sodium Iron Fluorine Cobalt Iodine Zinc Selenium Silicon Nickel Arsenic Boron Copper Manganese Molybdenum Chromium

15. Atoms Smallest “particle” of an element.

16. Molecule Smallest “particle” of a pure substances (element or compound) bonded together. –Combination of similar atoms (O 2 - element) –Combination of different atoms (H 2 O - 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 Elements – A large collection of atoms of a given type. Atoms of element A Atoms of element B Atoms of element A & B existing together Molecules made from element A & B through a process called bonding. Compound – a large collection of molecules made from atoms from element A & B B A

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

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 of isotopes 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 determines the size –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 Atomic Number Element Name Symbol or Abbrev. Atomic Mass Unit (AMU) 1 Hydrogen H 1.01

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 55

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. 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 Carbon and the amu

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

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 number changes (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 = 18 neutrons 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) Mass number Symbol Atomic number 14 O 8 15 O 8 16 O 8 Number of neutrons = (Mass # - Atomic #) ??? Number of protons & electrons ???

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