1. Lecture 3 Ch.2 Suggested HW: 15, 19, 23, 28, 38, 49, 52, 58 Lecture 3 Ch.2 Suggested HW: 15, 19, 23, 28, 38, 49, 52, 58 Atomic Structure, Isotopes, And Ions

2. Understanding the Nature of Atoms If you cut a piece of graphite from the tip of a pencil into smaller and smaller pieces, how far could you go? Could you divide it forever, or would the small pieces eventually become indivisible? You would eventually end up with atoms (translates to “indivisible” in greek) of pure carbon. You can not divide a carbon atom into smaller pieces and still have carbon

3. An atom is the smallest identifiable unit of an element The theory that all matter is composed of atoms grew out of two primary laws 1.Law of conservation of mass 2.Law of constant composition Matter

4. In a chemical reaction, atoms are not created or destroyed, but merely rearranged. In other words, the total mass of substances present before a reaction is equal to the total mass after. Law of Conservation of Mass

5. The relative amounts of each element in a given substance are always the same, regardless of how the substance was made. Law of Constant Composition

7. We have established that matter is comprised of atoms. But what are atoms made of? In the 1800’s, physicists conducted numerous experiments which revealed that the atom itself is made up of even smaller, more fundamental particles. The three types of sub-atomic particles that make up the atom are known as: electrons protons neutrons Atomic Structure

8. J.J. Thomson’s Cathode Ray Experiment (late 1800’s) No Electric Field With Electric Field Discovery of the Electron

9. Atoms are charge neutral. If electrons reside within an atom, then an equivalent number of positive charges must also exist, appropriately named protons. How do all these charges coexist? Thomson proposed the very first theoretical model of the atom, the so- called plum pudding model (PM) shown to the right. Electrons reside in a sea of uniform positive charge Plum Pudding Model

10. Ernest Rutherford sought to test the PM model using the gold foil experiment (below) A beam of positively charged α -particles were focused on a very thin sheet of gold Based on the PM model, this beam would pass right through the gold foil. In actuality, the beam was deflected at odd angles, with some α -particles bouncing directly back!! Protons and The Nucleus

11. True model of the atom is a dense, positively charged, proton-loaded nucleus surrounded by a sparse electron cloud ! The vast majority of an atom’s mass is contained within the nucleus. Nucleus Electron cloud α particles scattered, repulsed particles transmitted through cloud THE ATOM

12. Rutherford’s model was incomplete. For example, a hydrogen atom has one proton and one electron, but is only ¼th the mass of a helium atom which has two electrons and two protons. If all of the mass of an atom comes from its sub-atomic particles, how do we explain the unaccounted for mass? The answer is neutrons, particles that are equal in mass to protons, but with no electrical charge. Neutrons

13. ParticleRelative ChargeMass (amu) Charges shown in table are relative to the charge of a proton. A proton has an actual charge of 1.602 x 10 -19 Coulombs (C), an electron has a charge of -1.602 x 10 -19 C. Opposite charges attract! Like charges repel!! Subatomic Particles and Their Relative Masses & Charges

14. The number of protons in an atom is called the atomic number. An element is defined by its atomic number. (ex. only carbon has 6 protons) For a given element, the number of protons DOES NOT CHANGE In a neutral atom, the number of protons is equal to the number of electrons. 6 C Carbon 12.0107 Atomic # Elemental Symbols

15. 6 C Carbon 12.0107 Mass # Elemental Symbols The mass number of an element is the sum of its protons and neutrons. The mass #’s listed on the periodic table are average masses (the unit of atomic mass is the amu, or atomic mass unit). These averages are used because numerous variations of elements called isotopes exist in nature.

16. Isotopes are variations of elements with the same number of protons but different numbers of neutrons. mass number atomic number Three isotopes of carbon. The % abundances of each are shown. Isotopes

17. Transitional Page Avg. atomic mass is obtained using the % abundance and the isotope mass.

18. For the table below, fill in the numbers of protons, neutrons, and electrons for each isotope of carbon. Then, using the given abundances and isotope masses, calculate the average atomic mass of C. Does it match the reported value? Group Work

19. Boron has two isotopes, 10B and 11B. Using the given isotope masses, determine the % abundances of each isotope. Group Work

20. The nuclei of most naturally occurring isotopes are very stable, despite the massive repulsive forces that exist between the protons in the nucleus. A strong force of attraction between neutrons and protons known as the nuclear force counteracts this repulsion. As the number of protons increases, more neutrons are required to stabilize the atom. Stable nuclei up to atomic number 20 have equal numbers of protons and neutrons. For nuclei with atomic number above 20, the number of neutrons exceeds the protons to create a stable nucleus. Proton-Neutron Ratio and Radioactivity

21. Radioactive isotopes are unstable (high in energy). This instability is attributed to a neutron/proton ratio that is either too high or too low. To become stable, they spontaneously release particles or radiation to lower their energy. This release of energy is called radioactive decay. Proton-Neutron Ratio and Radioactivity

22. Property α βγ Charge2+2+ 1-1- 0 Mass6.64 x 10 -24 g9.11 x 10 -28 g0 Emitted Radiation Type High energy electron. Pure energy (Radiation) Penetrating PowerLow. Stopped by paper. Blocked by skin. Moderate. Stopped by aluminum foil. (10 α ) High. Can penetrate several inches of lead. (10000 α ) The three most common types of radioactive decay are alpha, beta, and gamma Radioactivity

24. Stable isotopes within the belt *too few protons ( β -decay) *too many protons ( α -decay) All elements beyond atomic number 83 are radioactive. Belt of Stability

25. Thus far, we’ve learned that each element has an exact number of protons. – For example, Hydrogen has only one proton. If you force a second proton onto the atom, you no longer have hydrogen… you now have Helium. We have also learned that atoms can have variable numbers of neutrons (isotopes). Next, we will discuss ions. Ions

26. Ions are electrically charged atoms, resulting from the gain or loss of electrons. Positively charged ions are called cations. You form cations when electrons are lost Negatively charged ions are called anions. You form anions when electrons are gained Ions

27. A cation is named by adding the word “ion” to the end of the element name Anions are named by adding the suffix –ide to the end of an element Lithium ion Sodium ion Magnesium ion Aluminum ion Chloride Sulfide Oxide Phosphide Ion Nomenclature

28. Fill in the missing information below ISOTOPE PNE ??131410 Group Work