CHEM 98 Atoms and the Periodic Table Lecture Notes PDF

CHEM 98 Atoms and the Periodic Table Reference Outline Fundamentals of General, Organic, and Biological Chemistry 8th Edition McMurry, Ballantine, Hoeger, Peterson Chapter 2 2.1 Atomic Theory and the Structure of Atoms 2.2 Elements and Atomic Number 2.3 I...

CHEM 98 Atoms and the Periodic Table Reference Outline Fundamentals of General, Organic, and Biological Chemistry 8th Edition McMurry, Ballantine, Hoeger, Peterson Chapter 2 2.1 Atomic Theory and the Structure of Atoms 2.2 Elements and Atomic Number 2.3 Isotopes and Atomic Weight 2.4 The Periodic Table 2.5 Some Characteristics of Different Groups 2.6 Electronic Structure of Atoms 2.7 Electron Configurations 2.8 Electron Configurations and the Periodic Table 2.9 Electron-Dot Symbols  2.1 Atomic Theory and the Structure of Atoms The Atom Smallest particle that an element can be divided into and still be identifiable From the Greek atomos, meaning “indivisible.” Looks cool, but not really the correct depiction of an atom  2.1 Atomic Theory Atomic Theory Chemistry is founded on four assumptions: 1. All matter is composed of atoms. 2. The atoms of a given element differ from the atoms of all other elements. 3. Chemical compounds consist of atoms combined in specific ratios. Only whole atoms can combine. 4. Chemical reactions change only the way atoms are combined in compounds.  2.1 Atomic Theory Atoms Composed of tiny subatomic particles called protons, neutrons, and electrons. Protons carry a positive electrical charge. Neutrons have a mass similar to that of a proton, but are electrically neutral. Electrons have a mass that is only 1/1,836 that of a proton and carry a negative electrical charge.  2.1 Atomic Theory Masses of atoms and subatomic particles are expressed on a relative mass scale. Atomic mass unit (amu): – Unit for describing the mass of an atom – Based on the mass of a carbon-12 atom  2.1 Atomic Theory Atoms are mostly empty space Nucleus – Dense core consisting of protons and neutrons Area Surrounding Nucleus – Electrons move rapidly through a 'large' volume of space.  2.1 Atomic Theory Relative size of a nucleus in an atom Similar to a pea in the middle of a baseball stadium.  2.1 Atomic Theory Opposite electrical charges attract each other – Negatively charged electrons held near positively charged nucleus. Like charges repel each other – Electrons try to get as far away from each other as possible.  2.2 Elements and Atomic Number Atomic number (Z) Number of protons in atoms of a given element. – E.G. Na has 11 protons, Mg has 12 protons Mass number (A) Sum of protons and neutrons in an atom.  2.2 Elements and Atomic Number In Other Words: Atomic number (Z) = # of protons Mass number (A) = # of protons and neutrons Therefore: 𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑁𝑒𝑢𝑡𝑟𝑜𝑛𝑠 = 𝑀𝑎𝑠𝑠 𝑁𝑢𝑚𝑏𝑒𝑟 𝐴 − 𝐴𝑡𝑜𝑚𝑖𝑐 𝑁𝑢𝑚𝑏𝑒𝑟 𝑍 Silver (Ag) – Mass Number (A): 109 – Atomic Number (Z): 47 – Number of Neutrons = 109 – 47 = 62 109 47 Ag  2.2 Elements and Atomic Number Atoms are neutral (no net charge) # of protons (+) = # of electrons (-) In neutral atoms – Atomic number (Z) (# of protons) also equals the number of electrons E.G. Neutral Carbon, Z = 6, has 6 protons and 6 electrons  Worked Example 2.1 A Phosphorus has the atomic number Z = 15. Z P How many protons, electrons, and neutrons are there in phosphorus atoms, which have a mass number A = 31? It’s recommended that you have a periodic table nearby (either electronically or on paper) to look at as we talk about trends in the periodic in this lecture  Worked Example 2.2 A An atom contains 28 protons and has A = 60. Z ? Give the number for electrons and neutrons in the atom, and identify the element. Ni  2.3 Isotopes and Atomic Weight Isotopes Atoms with identical atomic numbers (Z) – Same number of protons But different mass numbers (A) – Differ in their number of neutrons Chlorine-35 35 37 Cl Cl Chlorine-37 Z = 17 (protons) Z = 17 (protons) A = 35 A = 37 18 neutrons 17 20 neutrons 17  2.3 Isotopes and Atomic Weight Isotopes of Hydrogen Protium, deuterium, and tritium  2.3 Isotopes and Atomic Weight A specific isotope is represented by: – Mass number (A) as a superscript – Atomic number (Z) as a subscript – Both in front of the atomic symbol For example, the symbol for tritium is:  2.3 Isotopes and Atomic Weight The isotopes of most elements do not have distinctive names (unlike hydrogen). The mass number (A) is given after the name of the element. – E.G. Uranium-235 or U-235 (used in nuclear reactors) – Compared to U-238 (non-radioactive) Most naturally occurring elements are mixtures of isotopes. 2.3 Isotopes and Atomic Weight  2.3 Isotopes and Atomic Weight Atomic weight Weighted average mass of an element’s atoms. Takes isotopic abundance into account Need to know: – Individual masses of the naturally occurring isotopes – Relative percentage of each isotope Calculated as the sum of the masses of the individual isotopes for that element. Atomic weight = Σ[(isotopic abundance) × (isotopic mass)]  Worked Example 2.3 Gallium Metal with a very low melting point (29.76 °C) – Will melt in the palm of your hand – ‘Disappearing Spoon’ in hot water Two naturally occurring isotopes: – 60.4% is Ga-69 (mass = 68.9257 amu) – 39.6% is Ga-71 (mass = 70.9248 amu) Calculate the atomic weight for gallium  Worked Example 2.4 Identify element X in the symbol: Give its: – Atomic number – Mass number 194 X – Number of protons – Number of electrons – Number of neutrons. 78  2.4 The Periodic Table Periodic table: A tabular format listing all known elements Atomic symbol, name of the element, and atomic mass are given in each box that represents the element  2.4 The Periodic Table Dmitri Mendeleev Produced the forerunner of the modern periodic table Boxes for each element with the symbol, atomic number, and atomic weight. 2.4 The Periodic Table  2.4 The Periodic Table Elements can be classified by Physical Properties Chemical Behaviour  2.4 The Periodic Table Physical Properties: Metal: – Malleable element – Lustrous appearance – Good conductor of heat and electricity – Metals occur on the left side of the periodic table. Sodium Magnesium Metal Metal  2.4 The Periodic Table Nonmetal: – Element that is a poor conductor of heat and electricity – Nonmetals occur on the upper-right side of the periodic table. Sulphur  2.4 The Periodic Table Metalloid: – Elements whose properties are intermediate between those of a metal and a nonmetal – Metalloids are located in a zigzag band between the metals on the left and nonmetals on the upper-right side of the periodic table.  2.4 The Periodic Table Elements can also be classified by chemical behavior. Elements in the same vertical column (group) – Similar chemical properties – Can be categorized into the following three groups: – Main group elements – Transition metal elements – Inner transition metal elements  2.5 Some Characteristics of Different Groups Periodicity: A repeating rise-and-fall pattern A graph of atomic radius versus atomic number shows periodicity.  2.5 Some Characteristics of Different Groups Group 1A—Alkali metals Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr) Shiny, soft metals with low melting points React with water to form products that are highly alkaline Because of their high reactivity, alkali metals are never found in nature in a pure state. Group 1A 2.5 Some Characteristics of Different Groups Alkali metals in water https://www.youtube.com/watch?v=uixxJtJPVXk Group 1A  2.5 Some Characteristics of Different Groups Group 2A—Alkaline earth metals Beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra) Lustrous, silvery metals Less reactive than their neighbors in group 1A Never found in nature in a pure state Group 2A  2.5 Some Characteristics of Different Groups Group 7A—Halogens Fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At) Colourful and corrosive nonmetals Found in nature only in combination with other elements, such as with sodium in table salt (sodium chloride, NaCl) Group 7A  2.5 Some Characteristics of Different Groups Group 8A—Noble gases Helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn) Colourless gases Labeled the “noble” gases because of their lack of chemical reactivity Helium, neon, and argon don’t combine with any other elements. Krypton and xenon combine with very few. Group 8A  2.6 Electronic Structure of Atoms The properties of the elements are determined by the arrangement of electrons in their atoms. Quantum Mechanical Model – Developed by Erwin Schrödinger  2.6 Electronic Structure of Atoms Quantum Mechanical Model Helps us understand the electronic structure of atoms Electrons have both particle-like and wave-like properties Wave Function Equation describing the behaviour of electrons Electrons are not perfectly free to move. They are restricted to certain energy values, or quantized. Coincidence? Schrödinger Indiana Jones and the Raiders of the Lost Ark  2.6 Electronic Structure of Atoms A ramp is not quantized because it changes height continuously. Stairs are quantized because they change height by a fixed amount.  2.6 Electronic Structure of Atoms Wave Functions Provide an electron with an “address” within an atom Composed of shell, subshell, and orbital. Shell Grouping of electrons in an atom according to energy Shell number: 1 2 3 4 Electron capacity: 2 8 18 32 1st Shell 2nd Shell 3rd Shell Etc.  2.6 Electronic Structure of Atoms The farther a shell is from the nucleus The larger it is, the more electrons it can hold, and the higher the energies of those electrons. Within the shells Electrons are further grouped into subshells of four different types, identified as s, p, d, and f in order of increasing energy. Subshells Within each subshell Electrons are grouped into orbitals, regions of space within an atom where the specific electrons are most likely to be found.  2.6 Electronic Structure of Atoms s subshell has 1 orbital p has 3 orbitals d has 5 orbitals f has 7 orbitals. Each orbital holds two electrons, which differ in a property known as spin.  2.6 Electronic Structure of Atoms Different orbitals have different shapes and orientations S and p orbitals shown  2.6 Electronic Structure of Atoms Different orbitals have different shapes and orientations s, p, d and f orbitals shown  2.6 Electronic Structure of Atoms Shells Subshells  2.7 Electron Configurations Electron Configuration Exact arrangement of electrons in an atom’s shells and subshells Rules (3) for writing electron configurations – Determines the order in which shells and subshells are filled  2.7 Electron Configurations Schemes to help remember the order in which the orbitals are filled Use one or the other (or your own!)  2.7 Electron Configurations Rule 1: Electrons occupy the lowest energy orbitals available. This is complicated by “crossover” of energies above the 3p level. Rule 2: Each orbital can hold only two electrons, which must be of opposite spin.  2.7 Electron Configurations Rule 3: Two or more orbitals with the same energy are each half-filled by one electron before any one orbital is completely filled by the addition of the second electron. – The number of electrons in each subshell is indicated by a superscript. 2.7 Electron Configurations 2.7 Electron Configurations 2.7 Electron Configurations  2.7 Electron Configurations A graphic representation can be made by indicating each orbital as a line and each electron as an arrow. The head of the arrow indicates the electron spin. A shorthand using noble gas configurations is very useful for large atoms.  2.7 Electron Configurations These are the electron configurations for B – N in which the 2p shell begins to fill.  2.7 Electron Configurations These are the electron configurations for O – Ne in which the 2p shell is completed.  Worked Example 2.6 Show how the electron configuration of magnesium can be assigned.  2.8 Electron Configurations and the Periodic Table The periodic table can be divided into blocks of elements according to the subshell filled last.  2.8 Electron Configurations and the Periodic Table Elements in the same group of the periodic table have similar electron configurations in their valence shells. A valence shell is the outermost electron shell of an atom. A valence electron is an electron in the valence shell of an atom. 2.8 Electron Configurations and the Periodic Table  Worked Example 2.8 Write the electron configuration for the following elements Use both the complete and the shorthand notations Indicate which electrons are the valence electrons. (a) Na (b) Cl  2.9 Electron-Dot Symbols Electron-dot (Lewis) symbol An atomic symbol with dots placed around it to indicate the number of valence electrons  Worked Example 2.11 Write the electron-dot symbol for any element X in group 5A.