Atomic Structure, Isotopes, Ions and Moles Ch.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? 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

Matter 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 Law of constant composition Law of conservation of mass

Atomic Structure 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

Subatomic Particles and Their Relative Masses & Charges Relative Charge Mass (amu) Proton +1 1.007 Neutron 1.008 Electron -1 .000548

C Elemental Symbols 6 Atomic # Carbon 12.0107 Atomic # 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.

C Elemental Symbols 6 Mass # Carbon 12.0107 Mass # The mass number of an element is the sum of its protons and neutrons. The mass #’s listed on the periodic table are averages . The unit of atomic mass is the amu, or atomic mass unit, which is equal to 1/12 of the mass of a carbon atom having 6 neutrons. These averages are used because numerous variations of elements called isotopes exist in nature.

Isotopes Isotopes are variations of elements having different numbers of neutrons. Isotope symbols are shown below for the three isotopes of nitrogen with their % abundances in nature. The 35Cl and 37Cl isotopes have 18 and 20 neutrons, respectively. mass number 𝟏𝟕 𝟑𝟓 𝑪𝒍 (75.78%) For any sample containing chlorine atoms, you will have the mix of isotopes shown. atomic number 𝟏𝟕 𝟑𝟕 𝑪𝒍 (24.22%)

Transitional Page Avg. atomic mass is obtained using the % abundance and the isotope mass. 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠= 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑎𝑠𝑠 𝑥 (𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒)

Sum of abundances must be 1 !!!! Example Confirm the average mass of Cl shown on the periodic table. 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑚𝑎𝑠𝑠= 34.968 852 71 amu 0.7578 +(36.965 902 60)(0.2422) 𝒎𝒂𝒔𝒔 𝒐𝒇 𝟏𝟕 𝟑𝟓 𝑪𝒍 𝒊𝒏 𝒂𝒎𝒖 𝒎𝒂𝒔𝒔 𝒐𝒇 𝟏𝟕 𝟑𝟕 𝑪𝒍 𝒊𝒏 𝒂𝒎𝒖 𝒂𝒃𝒖𝒏𝒅𝒂𝒏𝒄𝒆 𝒐𝒇 𝟏𝟕 𝟑𝟓 𝑪𝒍 𝒊𝒏 𝒂𝒎𝒖 𝒂𝒃𝒖𝒏𝒅𝒂𝒏𝒄𝒆 𝒐𝒇 𝟏𝟕 𝟑𝟓 𝑪𝒍 𝒊𝒏 𝒂𝒎𝒖 Sum of abundances must be 1 !!!! 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑚𝑎𝑠𝑠=35.453 𝑎𝑚𝑢

Group Work Boron has two isotopes, 10B and 11B. Using the given isotope masses, determine the % abundances of each isotope. ISOTOPE % A Mass (amu) 𝟓 𝟏𝟎 𝑩 10.013 𝟓 𝟏𝟏 𝑩 11.009

Ions 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 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

Ion Nomenclature 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 𝑳𝒊 + Lithium ion Sodium ion Magnesium ion Aluminum ion 𝑪𝒍 − Chloride Sulfide Oxide Phosphide 𝑵𝒂 + 𝑺 𝟐− 𝑴𝒈 𝟐+ 𝑶 𝟐− 𝑨𝒍 𝟑+ 𝑷 𝟑−

Group Work Fill in the missing information below ISOTOPE P N E 16 32 𝑆 16 32 𝑆 2- ?? 13 14 10 ?? ?? 𝑃𝑡 4+ 95