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 2 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 each element has a unique Name and Symbol symbol either one or two letters one capital letter or one capital letter + one lower case 4 1
The Periodic Table of Elements 5 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 10-24 g ~ 1 amu volume of H atom = 2.1 x 10-25 cm 3 6 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 neutral as #proton (#p) = #electron (#e) Nucleus Proton Neutron Electron 7 2
Comparison among Proton, Electron, Neutron Subatomic Particle Mass g Mass amu Location in atom Charge Symb ol Proton 1.67 x 10-24 1 nucleus +1 p, p +, H + Electron 9 x 10-28 ~0 empty space -1 e, e - Neutron 1.67 x 10-24 1 nucleus 0 n, n 0 8 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. 9 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 10 3
Isotopes Atomic Number (Z) Number of protons Mass Number (A) Protons + Neutrons Abundance = relative amount found in a sample Example: Cl-35 (75%) vs. Cl-37 (25%) 11 Isotopic Symbol Cl-35 has a mass number = 35, 17 protons and 18 neutrons (35-17). The symbol for this isotope would be 35 Cl 17 Atomic Symbol A = mass number Z = atomic number #neutrons = A - Z A X Z 12 Example: How many protons, neutrons, and electrons in an atom of 238 U 92 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 13 4
Write Isotopic symbol for the following two isotopes: a. Hydrogen isotope w/ 1 neutron b. Uranium isotope w/ 143 neutrons (#p + #n) #p Element 2 H 1 235 92 U 14 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 15 Example: Find the atomic weight of the element chlorine Information: exact mass number Cl-35 (34.97 amu), Cl-37 (36.97 amu); isotopic abundance Cl-35 (75.78%), Cl-37 (24.22%) Atomic weight = sum of weighted atomic mass from all isotopes Mass due to Cl-35 = 34.97 amu 0.7578 Mass due to Cl-37 = 36.97 amu 0.2422 Atomic weight = 35.45 amu 16 5
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 17 The Modern Periodic Table Main Group = Representative Elements = A groups Transition Elements = B groups all metals Bottom rows = Inner Transition Elements = Rare Earth Elements metals really belong in Period 6 & 7 18 Main group vs. Transition metals, Inner transition metals = Metal = Metalloid IA IIA = Nonmetal IIIA VIIIA IIIB VIIB VIIIB IB IIB 6
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 20 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 21 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 22 7
= Alkali Metals = Alkaline Earth Metals = Noble Gases = Halogens = Lanthanides = Actinides = Transition Metals 23 = Transition Metals = Rare Earth Metals = Transuranium element U 24 Important Element - Hydrogen nonmetal colorless, diatomic gas H 2 very low melting point & density reacts with Nonmetals to form molecular compounds HCl is acidic gas H 2 O is a liquid reacts with Metals to form hydrides metal hydrides react with water to form H 2 Nickel-metal hydride (NiMH) used in rechargeable battery HX dissolves in water to form acids 25 8
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 26 Important Element - Carbon 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 27 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 H 2. 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 28 9
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 H 2, Be = none; Mg = steam; Ca, Sr, Ba = cold water beryllium magnesium calcium strontium barium 29 Group VIIA: Halogens nonmetals F 2 & Cl 2 gases; Br 2 liquid; I 2 solid all diatomic very reactive Cl 2, Br 2 react slowly with water Cl 2 + H 2 O HCl + HOCl (chlorine) react with metals to form ionic compounds HX all acids HF weak < HCl < HBr < HI fluorine chlorine bromine iodine 30 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 31 10
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% 32 Quantum-Mechanical Model: Quantum Numbers Three quantum numbers: quantize the energy Principal quantum number, n, specifies the main energy level for the orbital. Also called Shell Number. the higher n value, the higher energy of the electrons, the further away electrons are located from the nucleus 33 Quantum-Mechanical: Quantum Numbers Principal energy shell has one or more Subshells the number of subshells = the Principal quantum number n = 1, one subshell; n = 2, two subshells; n = 3, three subshells Subshell Quantum numbers: s, p, d, f each Subshell has orbitals with a particular shape the shape represents the probability map 90% probability of finding electron in that region 34 11
Shapes of Subshells s Orbital p Orbitals: p x, p y, p z d Orbitals 35 f orbitals Tro: Chemistry: A Molecular Approach, 2/e 36 36 Shapes of f orbitals: 4f orbitals (downloaded from public domain) The coloration corresponds to the sign of function. 37 12
Shells & Subshells 38 How does the 1s Subshell Differ from the 2s Subshell? 39 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: 2p x, 2p y and 2p z 40 13
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) 41 Energy 7s 6s 5s 4s 6p 5p 4p 3p 6 d 5d 4d 3d 5f 4f 3s 2p 2s 1s 42 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) 2. draw arrows through the diagonals, looping back to the next diagonal each time 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 43 14
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. 44 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 45 How electrons in an atom are filled into orbitals 1. How Electrons fill subshells with multiple orbitals 2. How Electrons fill subshells with higher n number first Energy level 7s 6d 7p 6p 5d 6s 5p 5s 4p 4s 3p 3s 2s 4d 2p 3d 5f 4f 1s 46 15
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: 2p 3 _ _ _ instead of 47 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 = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 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 = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 = [Kr]5s 1 48 Example: Ground State Orbital Diagram and Electron Configuration of Magnesium 1s 2 2s 2 2p 6 3s 2 = [Ne]3s 2 1s 2s 2p 3s 3p 49 16
Practice: Write Electron Configuration and Orbital Diagram for the following atoms at the Ground state Calcium Sulfur Potassium Phosphorus 50 Valence Electrons Definition: the electrons in all the subshells with the highest principal energy shell Example: electrons in bold Mg = [Ne]3s 2 O = [He]2s 2 2p 4 Br = [Ar]4s 2 3d 10 4p 5 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 51 Valence Electrons Rb = 37 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 the highest principal energy shell that contains electrons is the 5 th : 1 valence electron + 36 core electrons Kr = 36 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 the highest principal energy shell that contains electrons is the 4 th : 8 valence electrons + 28 core electrons 52 17
Electrons Configurations and the Periodic Table 53 Electron Configurations from the Periodic Table Example: Be 2s 2 B 2s 2 2p 1 C 2s 2 2p 2 N 2s 2 2p 3 O 2s 2 2p 4 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 2s 2 Ca 3s 2 Sr 4s 2 Ba 5s 2 VIIA: F 2s 2 2p 5 Cl 3s 2 3p 5 Br 4s 2 4p 5 I 5s 2 5p 5 54 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: ns 1 ; Group IIIA: ns 2 np 1 Group VIIA: ns 2 np 5 55 18
1 2 3 4 5 6 7 s 1 s 2 Electron Configuration & the Periodic Table p 1 p 2 p 3 p 4 p 5 s 2 d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 p 6 56 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 57 1 2 3 4 5 6 7 Electron Configuration from the Periodic Table 1A 2A 3s 2 3A 4A 5A 6A 7A P 3p 3 8A Ne P = [Ne]3s 2 3p 3 P has 5 valence electrons 58 19
1 2 3 4 5 6 7 Electron Configuration from the Periodic Table 1A 2A 4s 2 3d 10 3A 4A 5A 6A 7A As 4p 3 8A Ar As = [Ar]4s 2 3d 10 4p 3 As has 5 valence electrons 59 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 60 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 61 20
Everyone Wants to Be Like a Noble Gas! Alkali Metals (Group 1A) have one more electron than the previous noble gas, [NG]ns 1 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 + 62 Everyone Wants to Be Like a Noble Gas! Halogens (Group 7A) one fewer electron than the next noble gas: [NG]ns 2 np 5 Reactions with Metals: tend to gain an electron and attain the electron configuration of the next noble gas: [NG]ns 2 np 5 + 1e [NG]ns 2 np 6 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 63 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 64 21
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 Atom Atom s Electron Config Ion Ion s Electron Config Na [Ne]3s 1 Na + [Ne] Mg [Ne]3s 2 Mg 2+ [Ne] Al [Ne]3s 2 3p 1 Al 3+ [Ne] O [He]2s2p 4 O 2- [Ne] F [He]2s 2 2p 5 F - [Ne] 65 Trends in Atomic Size Down a group valence shell farther from nucleus effective nuclear charge fairly close Across a period (left to right) adding electrons to same valence shell effective nuclear charge increases valence shell held closer 66 Trends in Atomic Size 67 22
Metallic Character Metals 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 68 Trends in Metallic Character 69 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. 70 23
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 71 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 72 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 73 24