Atoms & the Periodic Table

Chapter Outline What is Atom? Chemical properties of Atoms: the Periodicity Isotopes Electrons in Atom: Quantum physics’ view Valence electrons and the Periodic Table Periodic trend: Atomic Radius, Metallic Character, Ionization Energy

Experiencing Atoms Atoms: incredibly small, yet compose everything atoms are the pieces of Elements properties of the atoms determine the properties of the elements

Within an Atom Atoms = (Protons + Neutrons) + Electrons The nucleus (Protons + Neutrons) is only about 10-13 cm in diameter yet with most of the mass of the atom The electrons move outside the nucleus with an average distance of about 10-8 cm The atom is electrically neutral : #proton (#p) = #electron (#e)

Comparison among Proton, Electron, Neutron

Elements each Element has a unique number of protons in its nucleus Atomic number: the number of Protons in the nucleus of an atom the elements are arranged on the Periodic Table in order of their atomic numbers Name and Symbol of an Element symbol either one or two letters one capital letter or one capital letter + one lower case

The Periodic Table of Elements

The Size of Atoms Atomic Mass Unit (amu): 1 amu = 1.66  10-24 g Hydrogen the smallest atom mass of H atom= 1.67 x 10-24g ~ 1 amu volume of H atom = 2.1 x 10-25cm3

Isotopes The same element could have atoms with different masses Examples: 2 isotopes of chlorine atoms in nature: one weighs about 35 amu (Cl-35); another weighs about 37 amu (Cl-37) Carbon-12 (C-12) is much more abundant than C-13.

Isotopes all isotopes of an element: chemically identical undergo the exact same chemical reactions the same number of protons different masses due to different numbers of neutrons. Example: C-14 atom has eight neutrons; C-12 atom has six neutrons. identified by their mass numbers protons + neutrons

Isotopes Atomic Number (Z) Mass Number (A) Number of protons Mass Number (A) Protons + Neutrons Abundance = relative amount found in a sample Example: Cl-35 (75%) vs. Cl-37 (25%)

Isotopic Symbol Cl-35 has a mass number = 35, 17 protons and 18 neutrons (35 - 17). The symbol for this isotope would be Cl 35 17 Atomic Symbol A = mass number Z = atomic number #neutrons = A - Z AX Z

Example: How many protons, neutrons, and electrons in an atom of Isotopic symbol  element atomic number #p  #e #n Mass number = Atomic number (# protons, or #p) + #neutrons U = uranium Atomic Number = 92 #p = atomic number = 92 #e = #p = 92 Mass Number = #p + #n 238 = 92 + #n 146 = #n #proton = 92 #neutron = 146 #electron = 92

Mass Number is Not the Same as Atomic Mass Atomic mass (or Atomic Weight) is an experimental number determined from all naturally occurring isotopes Mass number refers to the number of protons + neutrons in one isotope natural or man-made When given the relative abundance of all isotopes, we can find the Atomic mass

The Modern Periodic Table Elements with similar chemical and physical properties are in the same column columns are called Groups or Families designated by a number and letter at top rows are called Periods each period shows the pattern of properties repeated in the next period

The Modern Periodic Table Main Group = Representative Elements = ‘A’ groups Transition Metals = ‘B’ groups Aka Transition elements Inner Transition Elements = Bottom rows = Rare Earth Elements metals really belong in Period 6 & 7

Main group vs. Transition metals, Inner transition metals = Metalloid = Nonmetal IA VIIIA IIA IIIA IIIB VIIB VIIIB IB IIB

Metals: Physical vs. Chemical Properties solids at room temperature, except Hg reflective surface shiny conduct heat, electricity Malleable (can be shaped) Tend to Lose electrons and form Cations in reactions. Na  Na+ + e - about 75% of the elements are metals lower left on the table

Nonmetals: Physical vs. Chemical Properties Elements found in all 3 states poor conductors of heat or electricity solids are brittle Tend to gain electrons in reactions to become anions: Cl + e -  Cl- upper right on the table except H

Metalloids: between Metals and Nonmetals show some properties of metals and some of nonmetals also known as semiconductors Properties of Silicon shiny conducts electricity does not conduct heat well brittle

= Alkaline Earth Metals = Alkali Metals = Alkaline Earth Metals = Noble Gases = Halogens = Lanthanides = Actinides = Transition Metals add pictures of elements from text

= Transuranium element = Transition Metals = Rare Earth Metals = Transuranium element add pictures of elements from text U

Important Element - Hydrogen nonmetal colorless, diatomic gas H2 very low melting point & density reacts with Nonmetals to form molecular compounds HCl is acidic gas H2O is a liquid reacts with Metals to form hydrides metal hydrides react with water to form H2 Nickel-metal hydride (NiMH) used in rechargeable battery HX dissolves in water to form acids

Important Element - Carbon Three forms of pure carbon: Diamond: hardest substance in nature Graphite: soft and slippery solid Buckminsterfullerene: a molecule made of 60 (Images from public domain wikipedia.com) carbon atoms in a sphere

Carbon as backbone for Organic/Biochemical Molecules Carbon atoms capable of forming robust bonds with many other elements and themselves. Examples: Small molecules: Butane, Sugar, Fatty acid, Vitamins Big molecules (Polymers): Starch, Kevlar, Teflon, Protein, and DNA

Group IA: Alkali Metals Usually Hydrogen is included All metals: soft, low melting points Flame tests ® Li = red, Na = yellow, K = violet Chemical Property: Very reactive. React with water to form basic (alkaline) solutions and H2. releases a lot of heat Tend to form water soluble compounds, such as table salt and baking soda. colorless solutions lithium sodium potassium rubidium cesium

Group IIA: Alkali Earth Metals Physical properties: harder, higher melting, and denser than alkali metals flame tests ® Ca = red, Sr = red, Ba = yellow-green Chemical properties: reactive, but less than corresponding alkali metal form stable, insoluble oxides. oxides are basic = alkaline earth reactivity with water to form H2, ® Be = none; Mg = steam; Ca, Sr, Ba = cold water beryllium magnesium calcium strontium barium

Group VIIA: Halogens nonmetals F2 & Cl2 gases; Br2 liquid; I2 solid all diatomic very reactive Cl2, Br2 react slowly with water Cl2 + H2O ® HCl + HOCl (chlorine) react with metals to form ionic compounds HX all acids HF weak < HCl < HBr < HI fluorine chlorine bromine iodine

Group VIIIA: Noble Gases all gases at room temperature, very low melting and boiling points very unreactive, practically inert very hard to remove electron from or give an electron to

Atomic Orbitals Quantum Physicists including Schrödinger: Electrons move very fast around the nucleus Electrons show up with a particular probability at certain location of the atom Orbital: A region where the electrons show up a very high probability when it has a particular amount of energy generally set at 90 or 95%

Electron Shells Cl (17p+ & 17e-) Electron Shell: the main energy level for the orbital. Principal quantum number n = 1, 2, … For a chlorine atom, three shells of electrons: The innermost shell (n = 1, 2 electrons) has the lowest energy The outmost shell (n = 3, 7 electrons) has the highest energy 17 p+ 2e- 8e- 7e- Cl (17p+ & 17e-)

Each Shell have Subshells Each Electron Shell has one or more Subshells (s, p, d, f) the number of subshells = the Principal quantum number n n = 1, one subshell (1s); n = 2, two subshells (2s, 2p) n = 3, three subshells (3s, 3p, 3d) n = 4, three subshells (4s, 4p, 4d, 4f) each Subshell has orbitals with a particular shape the shape represents the probability map 90% probability of finding electron in that region

Shapes of Subshells s Orbital p Orbitals: px , py , pz d Orbitals

f orbitals Tro: Chemistry: A Molecular Approach, 2/e 34

Shapes of f orbitals: 4f orbitals (downloaded from public domain) The coloration corresponds to the sign of function.

Shells & Subshells

How does the 1s Subshell Differ from the 2s Subshell?

Subshells and Orbitals Among the subshells of a principal shell, slightly different energies for multielectron atoms, the subshells have different energies: s < p < d < f each subshell contains one or more Orbitals s : 1 orbital p : 3 orbitals d : 5 orbitals f : 7 orbitals within one subshell, different orbitals have the same energy. Example: 2px, 2py and 2pz

Electron Configurations Definition: The distribution of electrons into the various energy shells (n = 1,2,3,…) and subshells (s, p, d, f) in an atom in its ground state Each energy shell and subshell has a maximum number of electrons it can hold Subshell s = 2, p = 6, d = 10, f = 14 Shell n: 1 = 2e, 2 = 8e, 3 = 18e, 4 = 32e Electrons fill in the energy shells and subshells in order of energy, from low energy up Aufbau Principal (“Construction” in German)

6s 6p 6d 7s 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d Energy 2s 2p 1s

Order of Subshell Filling in Ground State Electron Configurations 1. Diagram putting each energy shell on a row and listing the subshells, (s, p, d, f), for that shell in order of energy, (left-to-right) 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 2. draw arrows through the diagonals, looping back to the next diagonal each time

Spinning Electron(s) in Orbital Experiments showed Electrons spin on an axis generating their own magnetic field Pauli Exclusion Principle each Orbital may have a maximum of 2 electrons, with opposite spin Two electrons sharing the same orbital must have Opposite spins so there magnetic fields will cancel analogous to two bar magnets in parallel: only opposite alignment could stabilize each other.

Orbital Diagrams often an orbital as a square the electrons in that orbital as arrows the direction of the arrow represents the spin of the electron unoccupied orbital orbital with 1 electron orbital with 2 electrons

Filling the Orbitals in a Subshell with Electrons Energy shells fill from lowest energy to high 1 → 2 → 3 → 4 Subshells fill from lowest energy to high s → p → d → f Orbitals of the same subshell have the same energy. Three 2p orbitals; Five 3d orbitals  Electrons prefer “spreading out” in orbitals of same subshell before they pair up in orbitals. Hund’s Rule Example: 2p3 _ _ _ instead of  ____

Electron Configuration of Atoms in their Ground State Electron configuration: a listing of the subshells in order of filling with the number of electrons in that subshell written as a superscript Kr = 36 electrons = 1s22s22p63s23p64s23d104p6 a shorthand way : use the symbol of the previous noble gas in [] for the inner electrons, then just write the last set Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1

Example: Ground State Orbital Diagram and Electron Configuration of Magnesium 1s22s22p63s2 = [Ne]3s2 1s 2s 2p 3s 3p 

Practice: Write Electron Configuration for the following atoms at the Ground state Calcium Sulfur Sodium Chlorine Important: you are required to be able to write electron configuration for first three rows of elements

Valence Electron vs. Core Electron Valence Electron: the electrons in all the subshells with the highest principal energy shell Example: electrons in bold Mg = [Ne]3s2 O = [He]2s22p4 Br = [Ar]4s23d104p5 Core electrons: electrons in lower energy shells Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the Number of Valence electrons

Valence Electrons Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 the highest principal energy shell that contains electrons is the 5th : 1 valence electron + 36 core electrons Kr = 36 electrons = 1s22s22p63s23p64s23d104p6 the highest principal energy shell that contains electrons is the 4th : 8 valence electrons + 28 core electrons

Electrons Configurations and the Periodic Table

Electron Configurations from the Periodic Table Example: Be 2s2 B 2s22p1 C 2s22p2 N 2s22p3 O 2s22p4 Elements in the same period (row) have Valence Electrons in the same principal energy shell. #Valence electrons increases by one from left to right Elements in the same group have the same #valence electron and they are same kind of subshell Example: IIA: Be 2s2 Ca 3s2 Sr 4s2 Ba 5s2 VIIA: F 2s22p5 Cl 3s23p5 Br 4s24p5 I 5s25p5

Electron Configuration & the Periodic Table Elements in the same Group have similar chemical and physical properties  their valence shell electron configuration is the same No. Valence electrons for the main group elements is the same as the Group Number Example: Group IA: ns1 ; Group IIIA: ns2np1 Group VIIA: ns2np5

Electron Configuration & the Periodic Table s1 s2 p1 p2 p3 p4 p5 s2 1 2 3 4 5 6 7 p6 d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

Electron Configuration from the Periodic Table Inner electron configuration = Noble gas of the preceding period Outer electron configuration: from the preceding Noble gas the next period (Subshells)  Element the valence energy shell = the period number the d block is always one energy shell below the period number and the f is two energy shells below

Electron configuration & Chemical Reactivity Chemical properties of the elements are largely determined by No. Valence electrons Why elements in groups? Since elements in the same column have the same #valence electrons, they show similar properties

Electron Configuration: Noble Gas Noble gases have 8 valence electrons except for He, which has only 2 electrons Noble gases are especially nonreactive He and Ne are practically inert  The reason: the electron configuration of the noble gases is especially stable

Everyone Wants to Be Like a Noble Gas! Alkali Metals (Group 1A) have one more electron than the previous noble gas, [NG]ns1 tend to lose their extra ONE electron, resulting in the same electron configuration as a noble gas forming a cation with a 1+ charge Na  Na+ Li  Li+

Everyone Wants to Be Like a Noble Gas! Halogens (Group 7A) one fewer electron than the next noble gas: [NG]ns2np5 Reactions with Metals: tend to gain an electron and attain the electron configuration of the next noble gas: [NG]ns2np5 + 1e  [NG]ns2np6 forming an anion with charge 1-: Cl  Cl- Reactions with Nonmetals: tend to share electrons so that each attains the electron configuration of a noble gas

Everyone Wants to Be Like a Noble Gas! Summary Alkali Metals as a group are the most reactive metals they react with many things and do so rapidly Halogens are the most reactive group of nonmetals one reason for their high reactivity: they are only ONE electron away from having a very stable electron configuration the same as a noble gas

Stable Electron Configuration And Ion Charge Metals:  Cations by losing enough electrons to get the same electron configuration as the previous noble gas Nonmetals:  Anions by gaining enough electrons to get the same electron configuration as the next noble gas

Trends in Atomic Size Increases down a group valence shell farther from nucleus effective nuclear charge fairly close Decreases across a period (left to right) adding electrons to same valence shell effective nuclear charge increases valence shell held closer

Trends in Atomic Size

Metallic Character Metals Nonmetals malleable & ductile shiny, lusterous, reflect light conduct heat and electricity most oxides basic and ionic form cations in solution lose electrons in reactions – oxidized Nonmetals brittle in solid state dull electrical and thermal insulators most oxides are acidic and molecular form anions and polyatomic anions gain electrons in reactions - reduced

Trends in Metallic Character

Electron Configuration Affects the Size of Atoms and Metallic Character: Within a Group Within the same Group, from top to bottom: As quantum number n increases for the valence electron(s) valence electron(s) further away from the nucleus Larger Atomic Radius weaker Coulombic force (electrostatic force) withholding valence electrons electrons easier to be lost Stronger metallic character

Be (4p+ & 4e-) Mg (12p+ & 12e-) Ca (20p+ & 20e-) Example: Group IIA

Electron Configuration Affects the Size of Atoms and Metallic Character: Over the Period Within the same Period (row), from left to right: Same quantum number n for the valence electron(s) As Nucleus has increasing number of protons (p+) Stronger Coulombic force (electrostatic force) withholding valence electrons Valence Electrons closer the nuclues Smaller Atomic Radius Valence electrons harder to be lost Weaker metallic character

Li (3p+ & 3e-) Be (4p+ & 4e-) B (5p+ & 5e-) O (8p+ & 8e-) Example: Period 2 2e- 1e- 3+ 2e- 4+ 2e- 3e- 5+ Li (3p+ & 3e-) Be (4p+ & 4e-) B (5p+ & 5e-) 6+ 2e- 4e- 8+ 2e- 6e- 10+ 2e- 8e- O (8p+ & 8e-) C (6p+ & 6e-) Ne (10p+ & 10e-)

Ionization Energy (IE) In an atom, electrons (“-” charge) are attracted to the nucleus (“+” charge). Energy is required to remove the electron from an atom. Na + energy  Na+ + e- Neutral atom IE Cation Higher IE corresponds to lower Metallic property.

Trends in Ionization Energy Decreases down a group valence shell farther from nucleus effective nuclear charge fairly close Increases across a period (left to right) adding electrons to same valence shell effective nuclear charge increases valence shell held closer

Practice: Rank elements K, Mg, S, F: A. increasing metallic character B. increasing atomic radii C. increasing ionization energy F, S, Mg, K K, Mg, S, F