Homework

Atomic Particles, Mass, Weight, and Isotopes

Chapter 2: 1-2 Homework

Textbook assignment: ReadKotz and Triechel, Chemistry and Chemical Reactivity Chapter 2: Sections 1-2.

Study Notes
• 2.1 Atomic structure When Dalton resurrected the ancient Greek concept of atomi to explain whole-number ratios of elements in compounds, and fixed-formula compound structure, he assumed, as Democritus had, that the elemental atoms were indivisible, and this "rule" persisted until observations of radiation and electricity convinced the Curies and Thompson that atoms themselves contained components that could be rearranged or ejected from inside. Rutherford's alpha-particle experiment offered further evidence that atoms are mostly space, and that the component stable particles (protons, neutrons, and electrons) are very small compared to the overall size of the atom.
• 2.1 (con't) Atomic number and mass Atomic number identifies the element, atomic mass identifies (with atomic number) the isotope. Atomic number is simply the number of protons in the atom. All carbon atoms, for instance, have 6 protons, regardless of the number of electrons or neutrons they may also have at any given instant. It is the number of protons that makes the atom a particular element. Atomic mass is calculated using atomic mass units, where an amu = 1/12 * the mass of a carbon-12 atom with six protons and six neutrons. Because the number of neutrons in atoms of the same element may vary, the average atomic mass of an sample can be a fractional value. We can have a carbon sample, where some atoms have atomic mass 12 (6 protons + 6 neutrons), and other carbon atoms with atomic mass 14 (6 protons + 8 neutrons) so the [total mass]/[number of atoms] will be around 12.11 amu. Number of protons and total number of "nucleons" (protons plus neutrons) identify isotopes.
• 2.2 Isotopes and atomic weight Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons. Because the number of neutrons differs, the mass differs. When mass is given in amu, the mass for a single isotope will be a whole number. In any naturally-occurring sample of a element with multiple isotopes, there will be some atoms of each. You need to be able to figure the relative abundances of different isotopes given the atomic number, the atomic weight, and the number of neutrons in each isotope. The atomic weight is an average based on experimental data: the masses of all the atoms of all the isotopes divided by the total number of atoms. You may discover differences in atomic weights or atomic mass if you look at periodic tables created decades apart. One of the reasons is that our definition of an atomic mass unit has changed from a basis on the mass of hydrogen to a basis as 1/12 of the mass of carbon-12. While this can be confusing, the key to chemistry is proportionality -- which doesn't change when units change. We use atomic weights when we attempt to analyze or predict the presence of an element in a sample.
• HISTORICAL NOTES: The evidence for atomic structure and periodic table relationships comes from several key experiments.
• John Dalton's experiments in the early 1800s showed that elements combined in whole-number ratios whether measured by mass or by volume, and that different ratios of the same elements had different properties: H2O (water) and H2O2 (peroxide) are distinct substances.
• J. J. Thomson's experiments with cathode ray tubes revealed that particles smaller than hydrogen, the smallest element, existed. Thomson's beta rays (later identified as electrons) proved the atom itself was divisible.
• The Curies' experiments with radiation not only showed again that atoms had smaller component particles, but that atoms could change from one element to another by losing or gaining components.
• Rutherford's experiments demonstrated that atoms were mostly space.
• Millikan's experiment allowed him to determine the charge-to-mass ratio on the electron.

Important Formulae and Notation

Concept Formulae or Notation Explanation
Element Representation $X Z A$ X: Element Symbol
A: Mass number (protons+neutrons)
Z: Atomic Number (protons)
Percent Abundance Ai: count of atoms of specific isotope in sample
Atot: Count of all atoms in sample as determined from mass and atomic weight
Atomic Weight Pi: Percent abundance of isotope i
Mi: Atomic mass of isotope i

Web Lecture

You have two weblectures for this section, since some historical background is useful in understanding the observations and experimental evidence that led to our current understanding of atomic structure.

Read the following weblectures before chat: Atoms and Isotopes AND A History of Chemistry before the Curies (Optional: some of this information has also been incorporated into other weblectures)

Videos for Chapter 2

Our textbook publisher has a video website at Thinkwell Video Lessons

• Review the historical evidence that led to the structure of the atom in Early Discoveries and the Atom.
• Review Millikan's Oil Drop experiment with Understanding Electrons.
• Review the setup and results for Rutherford's experiment Understanding the Nucleus, which determined that atoms were mostly empty space.

Use the simulator below to experiment with atomic structure for atoms of different elements.

• Click on the "Atom" option.
• Click on the plus signes to display the periodic table, net charge, mass number, and stability.
• Start building an atom with 1 proton and 1 electron.
• Where do you need to drag the proton? Where do you need to drag the electron? What is the mass number? What is the net charge? Which element is highlighted on the period table?
• Drag a neutron to the center of your atom. What happens to the element in the periodic table? Is the atom stable or unstable?
• Drag a second neutron to the center of your atom. What happens to the mass number? What happens to the net charge? What happens to the atom's stability?
• Remove the electron from your atom. What happens to the mass number? What happens to the net charge? What happens to the atom's stability?
• Repeat the above exercise starting with an atom with 6 protons, 6 neutrons, and six electrons.
• What happens if you add one or two neutrons to the nucleus (center) of the atom?
• What happens if you remove an electron from the atom?
• What happens if you add a proton to the center of the atom?
• Click on the "Symbol" option.
• Again, create an atom with one proton and one electron. What is the symbol for this atom? What is the atomic number (lower left)? What is the atomic mass (upper left)? What is the charge (upper right)?
• Add a neutron. What changes in the symbol?
• Remove the electron, What changes in the symbol?
• Create an atom with six protons, six neutrons, and six electrons.
• How many neutrons do you need to add before the atom becomes unstable?
• Vary the number of electrons. What happens to the symbol?
• Add a proton. What happens to the symbol?
• Play one of the games to test your understanding of atomic structure, charge, and symbols.

Chat Preparation Activities

• Essay question: The Moodle forum for the session will assign a specific study question for you to prepare for chat. You need to read this question and post your answer before chat starts for this session.
• Mastery Exercise: The Moodle Mastery exercise for the chapter will contain sections related to our chat topic. Try to complete these before the chat starts, so that you can ask questions.

Chapter Quiz

• There is no chapter quiz YET.

Lab Work

Please read through the chapters on equipping your home lab and safely storing chemicals in Illustrated Guide to Hom Chemistry Experiments, chapters 3 and 4. We will discuss lab safety in the AP/lab chat. You should be acquiring the chemicals for this year's labs and be ready to start developing and demonstrating lab skills and safe lab practices.