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Unit 3 – Atoms : The Building Blocks of Matter - Nuclear Chemistry

Learn about the basic structure of atoms and the properties of its subatomic particles. Understand the concepts of atomic number, mass number, isotopes, and molar mass.

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Unit 3 – Atoms : The Building Blocks of Matter - Nuclear Chemistry

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  1. Unit 3 – Atoms: The Building Blocks of Matter - Nuclear Chemistry Learning Objectives and Learning Target Packet -Text support Ch. 3 and 21

  2. Atomic Theory Continues… • We have come so far in our understanding of atoms. Centuries of researching and countless scientists devoting their lives to create the understanding of the atom today (textbook concepts). • However it is NOT over. The more we understand about atoms (how they work, their make up, etc…) the greater our ability to advance science and technology in all aspects of our lives (i.e. medicine).

  3. Check for Understanding • What is the charge of a proton? • Where would you find a proton in an atom? • What is the charge of an electron? • Where would you find an electron in an atom? • What is the charge of a neutron? • Where would you find a neutron in an atom? • How big is an electron compared to a proton? • How big is a neutron compared to a proton?

  4. The Atom • The smallest particle of an element that retains the chemical properties of that element. • Consists of two regions: • Nucleus • 1. Small region located at the center of an atom • 2. Made up of at least one positively charged particle (proton) • 3. Made up of usually one or more neutral particles (neutrons) • Region surrounding the nucleus – Electron Cloud • 1. Very large compared to size of nucleus • 2. Contains the negatively charged particles (electrons)

  5. Refresh: What are Atoms? • Atoms are tiny particles that determine properties of all matter. • Atoms are the building blocks for molecules. • Atoms form elements. • Element: A substance that cannot be broken down into simpler substances by chemical means.

  6. Parts of an Atom • Proton: A subatomic particle that has a positive charge and is found in the nucleus of the atom. • Neutron: A subatomic particle that has NO charge and is found in the nucleus of the atom. • Electron: A subatomic particle that has a negative charge and moves around the outside of the nucleus.

  7. Label the Atom

  8. Subatomic Particles

  9. Electron Orbital -Electrons orbit the nucleus in orbital clouds. -Electrons with different amounts of energy exist in different energy levels.

  10. The Electron Cloud Model

  11. Electrons in each energy level • Each energy level can hold a limited number of electrons. • The lowest energy level is the smallest and the closest to the nucleus.

  12. Atomic Number • The atomicnumber of an element is the number of protons in the nucleus of an atom of that element.

  13. -2 Charge of Atoms • Atoms are not charged even though they have particles that contain charges. • Atoms are neutral because they have EQUAL numbers of protons and electrons. Ex: Helium Atom Charge of 2 protons: Charge of 2 neutrons: Charge of 2 electrons: Total charge of He atom: +2 0 0

  14. Mass Number • The sum of the protons and neutrons in the nucleus is the mass number of that particular atom. (mass # = p + n)

  15. Atomic Number vs. Mass Number • Atomic Number: Equal to the number of protons in the nucleus of the atom. (number of electrons = the number of protons) • Mass Number: Equal to the number of protons AND neutrons in an atom’s nucleus. • Average Atomic Weight (below symbol on PT): • Weighted average of the atomic masses of the naturally occurring isotopes of an element • We will round this number to the nearest hundredth (TWO decimal places) • Example Oxygen’s average atomic mass is 15.9994 = 16.00

  16. Application- TOGETHER Iodine (I) Iron (Fe) Atomic number ____ Atomic Number ____ Atomic Mass ____ Atomic Mass ____ Number of Protons ____ Number of Protons____ Number of Neutrons ____ Number of Neutrons____ Number of electrons ____ Number of Electrons ____

  17. Application- ON YOUR OWN Nickel (Ni) Radon (Rn) Atomic number ____ Atomic Number ____ Atomic Mass ____ Atomic Mass ____ Number of Protons ____ Number of Protons____ Number of Neutrons ____ Number of Neutrons____ Number of electrons ____ Number of Electrons ____

  18. Isotopes • Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number.

  19. Isotope Examples • Carbon – 12 • 6 protons • 6 electrons • 6 neutrons • Carbon – 13 • 6 protons • 6 electrons • 7 neutrons • Carbon 14 • 6 protons • 6 electrons • 8 neutrons

  20. 6,7 Mass Numbers Li Symbol 3 Atomic Number Lithium Element Name Example

  21. 107 Ag 47 Silver One More Example

  22. Isotope Tables • Breaking it down… • To find the symbol – determine the atomic number of the element. This is the number of protons • To find the protons- determine the atomic number of the element. • To find the electrons – equal to the number of protons of a neutral atom • To find neutrons: Mass Number – Atomic Number = Number of Neutrons • To find Mass Number: Atomic Number + Number of Neutrons = Mass Number

  23. Isotopes

  24. Molar Mass • Molar Mass: mass in grams of one mole of a pure substance. • Units = g/mol • To calculate molar mass you must find the atomic mass (Units = g). Ex: Find the atomic mass of Al. Ex: Find the atomic mass of O2. Ex: Find the atomic mass of CH4. * To change atomic mass to molar mass just put the answer (g) over moles.

  25. Practice

  26. Warm-up • Identify the element: • 65 neutrons • 48 protons • 48 electrons • How many protons, neutrons, electrons in the following isotopes • Mn-56 S-31 • Determine the molar mass of the following • H2O • Na2CO3

  27. The Mole

  28. Mole Music Video • http://www.youtube.com/watch?v=oIkC7SRqXP0

  29. The Mole &Avogadro’s Number NA = 6.02x1023 http://ed.ted.com/lessons/daniel-dulek-how-big-is-a-mole-not-the-animal-the-other-one

  30. Counting Particles

  31. One mole of Carbon Scientists Counting Unit • Since atoms are so small, it is hard to count individual atoms. • To solve this problem, chemists count by moles. • Moles: SI unit for measuring the amount of a substance.

  32. Mole Continued • One mole of anything contains 6.02 x 1023 “particles” • Particles can be atoms, ions, molecules, electrons, formula units, etc. Avogadro’s Number: 6.02x1023 = 1 mole 602,200,000,000,000,000,000,000

  33. How Big is Avogadro’s Number? • One mole of sheets of paper stacked one on top of the other would reach beyond the solar system. • One mole of basketballs could create a new planet the size of the earth. • One mole of rice grains would cover the land masses of Earth in a depth of 75 meters. • If you had a mole of pennies and gave away 1 million dollars of it (100,000,000 pennies) a day to everyone in the world it would take you more than 3000 years to distribute all your money.

  34. Conversions with Avogadro’s Number • Start with what you know. • Use the following conversion factor: 6.02 x 1023particles 1 mole • Cancel units. • Solve (preform the math).

  35. “Mole Map”

  36. Practice Conversions: Moles to Particles • Determine how many particles of sucrose are in 3.50 moles of sucrose. • Determine the number of atoms in 2.50 mol of Zn.

  37. Practice Conversions: Moles to Particles 3. Given 3.25 mol AgNO3, determine the number of formula units. 4. Calculate the number of molecules in 1.15 mol of water.

  38. Warm-Up Calculate the number of molecules in 1.15 mol of water.

  39. “Mole Map”

  40. Practice Conversions: Particles to Moles • How many moles are in 5.75x1024 atoms Al?

  41. Practice Conversions: Particles to Moles • How many ATOMS are in 3.75x1024 molecules of CO2? • How many moles are in 2.50x1020 atoms of Fe?

  42. Practice, Practice, Practice • Mole Conversion Worksheet

  43. Check for Understanding • 1.5 moles Cu = ________ atoms Cu • 8.51 x 1023 S = ________ moles S Solutions: • 9.0 x 1023 • 1.41

  44. Molar Mass and Conversions Adding Mass-Mole and Mole-Mass Finally, put it all together and what do you get?!

  45. Molar Mass • Molar Mass: mass in grams of one mole of a pure substance. • Units = g/mol • To calculate molar mass you must find the atomic mass (Units = g). Ex: Find the atomic mass of Al. Ex: Find the atomic mass of O2. Ex: Find the atomic mass of CH4. * To change atomic mass to molar mass just put the answer (g) over moles.

  46. Mole-Mass Conversions Conversion Factor: Ex: While working in a Chemistry lab, Gary needs 3.00 moles of Mn for a chemical reaction. How much Mn does Gary need to mass? Moles 1 Mass (g) P. Table

  47. Practice with Mole-Mass Conversions • Determine the mass in grams of each of the following: • 3.57 mol Al • 42.6 mol Si • 3.45 mol Co • Determine the moles in each of the following: 4. 25.5 g Ag 5. 125 g Zn 6. 1.45 kg Fe

  48. Mass-Atoms Conversions • This involves two conversions. • Conversion #1: Mass to moles • Conversion #2: Moles to particles • Ex: How many atoms of gold are in a pure gold nugget having a mass of 25.0 g? Moles 1 “Particles” 6.02x1023 Mass (g)

  49. Practice Mass-Atom Conversions How many atoms are in the following samples? • 55.2 g Li • 0.230 g Pb • 45.6 g Si

  50. Reverse It…Atoms-Mass Conversions • This involves two conversions. • Conversion #1: particles to moles • Conversion #2: moles to grams. Moles 1 “Particles” 6.02x1023 Mass (g)

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