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### Mec chapter 2

1. 1. Copyright© The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 2 The Structure of the Atom and the Periodic TableDennistonToppingCaret7th Edition
2. 2. 2.1 Composition of the Atom • Atom - the basic structural unit of an element • The smallest unit of an element that retains the chemical properties of that element
3. 3. 2.1 Composition of the Atom Electrons, Protons, and Neutrons • Atoms consist of three primary particles • electrons • protons • neutrons • Nucleus - small, dense, positively charged region in the center of the atom - protons - positively charged particles - neutrons - uncharged particles
4. 4. 2.1 Composition of the Atom Characteristics of Atomic Particles • Electrons are negatively charged particles located outside of the nucleus of an atom • Protons and electrons have charges that are equal in magnitude but opposite in sign • A neutral atom that has no electrical charge has the same number of protons and electrons • Electrons move very rapidly in a relatively large volume of space while the nucleus is small and dense
5. 5. 2.1 Composition of the Atom Symbolic Representation of an Element Charge of particle Mass A C Z X Atomic Symbol of number the atom • Atomic number (Z) - the number of protons in the atom • Mass number (A) - sum of the number of protons and neutrons
6. 6. 2.1 Composition of the Atom Atomic Calculations number of protons + number of neutrons = mass number number of neutrons = mass number - number of protons number of protons = number of electrons IF positive and negative charges cancel, the atom charge = 0
7. 7. 2.1 Composition of the Atom
8. 8. 2.1 Composition of the Atom Atomic Composition Calculations Calculate the number of protons, neutrons, and electrons in each of the following: 11 5 B 55 26 Fe
9. 9. 2.1 Composition of the Atom Isotopes • Isotopes - atoms of the same element having different masses – contain same number of protons 4 – contain different numbers of neutrons Isotopes of Hydrogen Hydrogen Deuterium Tritium (Hydrogen - 1) (Hydrogen - 2) (Hydrogen - 3)
10. 10. 2.1 Composition of the Atom Isotopic Calculations • Isotopes of the same element have identical chemical properties • Some isotopes are radioactive • Find chlorine on the periodic table • What is the atomic number of chlorine? 17 • What is the mass given? 35.45 • This is not the mass number of an isotope
11. 11. 2.1 Composition of the Atom Atomic Mass • What is this number: 35.34? • The atomic mass - the weighted average of the masses of all the isotopes that make up chlorine • Chlorine consists of chlorine-35 and chlorine-37 in a 3:1 ratio • Weighted average is an average corrected by the relative amounts of each isotope present in nature
12. 12. 2.1 Composition of the Atom Atomic Mass Calculation Calculate the atomic mass of naturally occurring chlorine if 75.77% of chlorine atoms are chlorine-35 and 24.23% of chlorine atoms are chlorine-37 Step 1: convert the percentage to a decimal fraction: 0.7577 chlorine-35 0.2423 chlorine-37
13. 13. 2.1 Composition of the Atom Step 2: multiply the decimal fraction by the mass of that isotope to obtain the isotope contribution to the atomic mass: For chlorine-35: 0.7577 x 35.00 amu = 26.52 amu For chlorine-37 0.2423 x 37.00 amu = 8.965 amu Step 3: sum these partial weights to get the weighted average atomic mass of chlorine: 26.52 amu + 8.965 amu = 35.49 amu
14. 14. 2.1 Composition of the Atom Atomic Mass Determination • Nitrogen consists of two naturally occurring isotopes – 99.63% nitrogen-14 with a mass of 14.003 amu – 0.37% nitrogen-15 with a mass of 15.000 amu • What is the atomic mass of nitrogen?
15. 15. 2.1 Composition of the Atom Ions and Charges • Ions - electrically charged particles that result from a gain or loss of one or more electrons by the parent atom • Cation - positively charged – results from the loss of electrons – 23Na  23Na+ + 1e- • Anion - negatively charged – results from the gain of electrons – 19F + 1e-  19F-
16. 16. 2.1 Composition of the Atom Calculating Subatomic Particles in Ions • How many protons, neutrons, and electrons are in the following ions? 39 + 19 K 32 2- 16 S 24 2+ 12 Mg
17. 17. 2.2 Development of Atomic Theory• Dalton’s Atomic Theory - the first experimentally based theory of atomic structure of the atom
18. 18. Postulates of Dalton’s Atomic Theory2.2 Development of 1. All matter consists of tiny particles Atomic Theory called atoms 2. An atom cannot be created, divided, destroyed, or converted to any other type of atom 3. Atoms of a particular element have identical properties
19. 19. 2.2 Development of 4. Atoms of different elements have different properties Atomic Theory 5. Atoms of different elements combine in simple whole-number ratios to produce compounds (stable aggregates of atoms) 6. Chemical change involves joining, separating, or rearranging atoms Postulates 1, 4, 5, and 6 are still regarded as true.
20. 20. Subatomic Particles: Electrons, Protons, and Neutrons2.2 Development of • Electrons were the first subatomic Atomic Theory particles to be discovered using the cathode ray tube. Indicated that the particles were negatively charged.
21. 21. Evidence for Protons and2.2 Development of Neutrons Atomic Theory • Protons were the next particle to be discovered, by Goldstein – Protons have the same size charge but opposite in sign – A proton is 1,837 times as heavy as an electron • Neutrons – Postulated to exist in 1920’s but not demonstrated to exist until 1932 – Almost the same mass as the proton
22. 22. 2.4 The Periodic Law and the Periodic Table• Dmitri Mendeleev and Lothar Meyer - two scientists working independently developed the precursor to our modern periodic table• They noticed that as you list elements in order of atomic mass, there is a distinct regular variation of their properties• Periodic law - the physical and chemical properties of the elements are periodic functions of their atomic numbers
23. 23. 2.4 The Periodic Lawand the Periodic Table Classification of the Elements
24. 24. 2.4 The Periodic Lawand the Periodic Table Important Biological Elements
25. 25. Parts of the Periodic Tableand the Periodic Table2.4 The Periodic Law • Period - a horizontal row of elements in the periodic table. They contain 2, 8, 8, 18, 18, and 32 elements • Group - also called families, and are columns of elements in the periodic table. • Elements in a particular group or family share many similarities, as in a human family.
26. 26. Families of the Periodic Tableand the Periodic Table2.4 The Periodic Law • Representative elements - Group A elements • Transition elements - Group B elements • Alkali metals - Group IA • Alkaline earth metals - group IIA • Halogens - group VIIA • Noble gases - group VIIIA
27. 27. Category Classification ofand the Periodic Table Elements2.4 The Periodic Law • Metals - elements that tend to lose electrons during chemical change, forming positive ions • Nonmetals - a substance whose atoms tend to gain electrons during chemical change, forming negative ions • Metalloids - have properties intermediate between metals and nonmetals
28. 28. Classification of Elementsand the Periodic Table Metals2.4 The Periodic Law • Metals: – A substance whose atoms tend to lose electrons during chemical change – Elements found primarily in the left 2/3 of the periodic table • Properties: – High thermal and electrical conductivities – High malleability and ductility – Metallic luster – Solid at room temperature
29. 29. Classification of Elementsand the Periodic Table Nonmetals2.4 The Periodic Law • Nonmetals: – A substance whose atoms may gain electrons, forming negative ions – Elements found in the right 1/3 of the periodic table • Properties: – Brittle – Powdery solids or gases – Opposite of metal properties
30. 30. Classification of Elementsand the Periodic Table Metalloids2.4 The Periodic Law • Metalloids: – Elements that form a narrow diagonal band in the periodic table between metals and nonmetals • Properties are somewhat between those of metals and nonmetals • Also called semimetals
31. 31. Atomic Number and Atomic Massand the Periodic Table2.4 The Periodic Law • Atomic Number: – The number of protons in the nucleus of an atom of an element – Nuclear charge or positive charge from the nucleus • Most periodic tables give the element symbol, atomic number, and atomic mass
32. 32. Element Information in theand the Periodic Table Periodic Table2.4 The Periodic Law 20 atomic number Ca symbol Calcium name 40.08 atomic mass
33. 33. Using the Periodic Tableand the Periodic Table2.4 The Periodic Law • Identify the group and period to which each of the following belongs: a. P b. Cr c. Element 30 • How many elements are found in period 6? • How many elements are in group VA?
34. 34. 2.5 Electron Arrangement and the Periodic Table• The electron arrangement is the primary factor in understanding how atoms join together to form compounds• Electron configuration - describes the arrangement of electrons in atoms• Valence electrons - outermost electrons – The electrons involved in chemical bonding
35. 35. 2.5 Electron Arrangement and the Periodic Table Valence Electrons • The number of valence electrons is the group number for the representative elements • The period number gives the energy level (n) of the valence shell for all elements
36. 36. 2.5 Electron Arrangement Valence Electrons and Energy and the Periodic Table Level • How many valence electrons does Fluorine have? – 7 valence electrons • What is the energy level of these electrons? – Energy level is n = 2
37. 37. 2.5 Electron Arrangement and the Periodic Table Energy Level Electron Arrangement by
38. 38. 2.5 Electron Arrangement Valence Electrons - Detail and the Periodic Table • What is the total number of electrons in fluorine? – Atomic number = 9 – 9 protons and 9 electrons • 7 electrons in the valence shell, (n = 2 energy level), so where are the other two electrons? – In n = 1 energy level – Level n = 1 holds only two electrons
39. 39. Determining Electron Arrangement2.5 Electron Arrangement and the Periodic Table List the total number of electrons, total number of valence electrons, and energy level of the valence electrons for silicon. 1. Find silicon in the periodic table • Group IVA • Period 3 • Atomic number = 14 1. Atomic number = number of electrons in an atom • Silicon has 14 electrons
40. 40. Determining Electron Arrangement #22.5 Electron Arrangement and the Periodic Table List the total number of electrons, total number of valence electrons, and energy level of the valence electrons for silicon. 3. As silicon is in Group IV, only 4 of its 14 electrons are valence electrons • Group IVA = number of valence electrons 3. Energy levels: • n = 1 holds 2 electrons • n = 2 holds 8 electrons (total of 10) • n = 3 holds remaining 4 electrons (total = 14)
41. 41. Determining Electron Arrangement2.5 Electron Arrangement and the Periodic Table Practice List the total number of electrons, total number of valence electrons, and energy level of the valence electrons for: • Na • Mg • S • Cl • Ar
42. 42. 2.5 Electron Arrangement Energy Levels and Subshells and the Periodic Table PRINCIPAL ENERGY LEVELS • n = 1, 2, 3, … • The larger the value of n, the higher the energy level and the farther away from the nucleus the electrons are • The number of sublevels in a principal energy level is equal to n – in n = 1, there is one sublevel – in n = 2, there are two sublevels
43. 43. 2.5 Electron Arrangement Principal Energy Levels and the Periodic Table • The electron capacity of a principal energy level (or total electrons it can hold) is 2(n)2 – n = 1 can hold 2(1)2 = 2 electrons – n = 2 can hold 2(2)2 = 8 electrons • How many electrons can be in the n = 3 level? – 2(3)2 = 18 • Compare the formula with periodic table…..
44. 44. n = 1, 2(1)2 = 2 n = 2, 2(2)2 = 8n = 3, 2(3)2 = 18 n = 4, 2(4)2 = 32
45. 45. 2.5 Electron Arrangement Sublevels and the Periodic Table • Sublevel: a set of energy-equal orbitals within a principal energy level • Subshells increase in energy: s<p<d<f • Electrons in 3d subshell have more energy than electrons in the 3p subshell • Specify both the principal energy level and a subshell when describing the location of an electron
46. 46. 2.5 Electron Arrangement Sublevels in Each Energy Level and the Periodic Table Principle energy Possible level (n) subshells 1 1s 2 2s, 2p 3 3s, 3p, 3d 4 4s, 4p, 4d, 4f
47. 47. 2.5 Electron Arrangement Orbitals and the Periodic Table • Orbital - a specific region of a sublevel containing a maximum of two electrons • Orbitals are named by their sublevel and principal energy level – 1s, 2s, 3s, 2p, etc. • Each type of orbital has a characteristic shape – s is spherically symmetrical – p has a shape much like a dumbbell
48. 48. 2.5 Electron Arrangement Orbital Shapes and the Periodic Table • s is spherically symmetrical • Each p has a shape much like a dumbbell, differing in the direction extending into space
49. 49. Number of2.5 Electron Arrangement Subshell orbitals and the Periodic Table s 1 p 3 d 5 f 7 •How many electrons can be in the 4d •10
50. 50. 2.5 Electron Arrangement Quantum Mechanical Model and the Periodic Table Shell 4 • Each orbital within a sublevel contains a 4f •• •• •• •• •• •• •• maximum of 2 Increasing Energy electrons 4d •• •• •• •• •• • Energy increases as n, shell number Sublevel increases, but ALSO 4p •• •• •• increases as you move from s to p to d to f Orbital sublevels 4s •• Electron
51. 51. 2.5 Electron Arrangement Electron Spin and the Periodic Table • Electron configuration - the arrangement of electrons in atomic orbitals • Aufbau principle - or building up principle helps determine the electron configuration – Electrons fill the lowest-energy orbital that is available first – Remember s<p<d<f in energy – When the orbital contains two electrons, the electrons are said to be paired
52. 52. 2.5 Electron Arrangement and the Periodic Table Electron Filling Order
53. 53. 2.5 Electron Arrangement Rules for Writing Electron and the Periodic Table Configurations • Obtain the total number of electrons in the atom from the atomic number • Electrons in atoms occupy the lowest energy orbitals that are available – 1s first • Each principal energy level, n contains only n sublevels • Each sublevel is composed of orbitals • No more than 2 electrons in any orbital • Maximum number of electrons in any principal energy level is 2(n)2
54. 54. 2.5 Electron Arrangement Electron Distribution and the Periodic Table • This table lists the number of electrons in each shell for the first 20 elements • Note that 3rd shell stops filling at 8 electrons even though it could hold more
55. 55. 2.5 Electron Arrangement and the Periodic Table Orbital Energy-level Diagram
56. 56. Writing Electron Configurations2.5 Electron Arrangement and the Periodic Table • H • Li – Hydrogen has – Lithium has 3 only 1 electron electrons – It is in the – First two have lowest energy configuration level & lowest of Helium – 1s2 orbital – 3rd is in the – Indicate orbital of number of lowest energy electrons with a in n=2 superscript – 1s2 2s1 – 1s1
57. 57. 2.5 Electron Arrangement Electron Configuration Examples and the Periodic Table • Give the complete electron configuration of each element – Be –N – Na – Cl – Ag
58. 58. 2.5 Electron Arrangement and the Periodic Table The Shell Model and Chemical Properties • As we explore the model placing electrons in shells, we will see that the pattern which emerges from this placement correlates well with a pattern for various chemical properties • We will see that all elements in a group have the same number of electrons in their outermost (or valence) shell
59. 59. 2.5 Electron Arrangement Groups Have Similar Chemical and the Periodic Table Properties and Appearances • Examples of different elements that have similar properties and are all in group VA – Nitrogen – Phosphorus – Arsenic – Antimony – Bismuth
60. 60. 2.5 Electron Arrangement Shorthand Electron and the Periodic Table Configurations • Uses noble gas symbols to represent the inner shell and the outer shell or valance shell is written after • Aluminum- full electron configuration is: 1s22s22p63s23p1 What noble gas configuration is this? •Neon •Configuration is written: [Ne]3s23p1
61. 61. 2.5 Electron Arrangement and the Periodic Table • Remember: – How many subshells are in each principle energy level? – There are n subshells in the n principle energy level. – How many orbitals are in each subshell? – s has 1, p has 3, d has 5, and f has 7 – How many electrons fit in each orbital? – 2
62. 62. 2.5 Electron Arrangement Shorthand Electron and the Periodic Table Configuration Examples • N • S • Ti • Sn
63. 63. 2.5 Electron Arrangement Classification of Elements and the Periodic Table According to the Type of Subshells Being Filled Use this breakdown of the Periodic Table and you can write the configuration of any element.
64. 64. 2.5 Electron Arrangement and the Periodic Table Classification of Elements – by Group • Representative element: An element in which the distinguishing electron is found in an s or p subshell • Distinguishing electron: The last or highest- energy electron found in an element • Transition element: An element in which the distinguishing electron is found in a d subshell • Inner-transition element: An element in which the distinguishing electron is found in a f subshell
65. 65. 2.6 The Octet Rule• The noble gases are extremely stable – Called inert as they don’t readily bond to other elements• The stability is due to a full complement of valence electrons in the outermost s and p sublevels: – 2 electrons in the 1s of Helium – the s and p subshells are full in the outermost shell of the other noble gases (eight electrons)
66. 66. Octet of Electrons2.6 The Octet Rule • Elements in families other than the noble gases are more reactive – Strive to achieve a more stable electron configuration – Change the number of electrons in the atom to result in full s and p sublevels • Stable electron configuration is called the “noble gas” configuration
67. 67. 2.6 The Octet Rule The Octet Rule • Octet rule - elements usually react in such a way as to attain the electron configuration of the noble gas closest to them in the periodic table – Elements on the right side of the table move right to the next noble gas – Elements on the left side move “backwards” to the noble gas of the previous row • Atoms will gain, lose or share electrons in chemical reactions to attain this more stable energy state
68. 68. 2.6 The Octet Rule Ion Formation and the Octet Rule • Metallic elements tend to form positively charged ions called cations • Metals tend to lose all their valence electrons to obtain a configuration of the noble gas Na Na+ + e- Sodium atom Sodium ion 11e-, 1 valence e- 10e- [Ne]3s1 [Ne]
69. 69. 2.6 The Octet Rule Ion Formation and the Octet Rule • All atoms of a group lose the same number of electrons • Resulting ion has the same number of electrons as the nearest (previous) noble gas atom Al Al3+ + 3e- Aluminum atom Aluminum ion 13e-, 3 valence e- 10e- [Ne]3s23p1 [Ne]
70. 70. Isoelectronic • Isoelectronic - atoms of different elements having2.6 The Octet Rule the same electron configuration (same number of electrons) • Nonmetallic elements, located on the right side of the periodic table, tend to form negatively charged ions called anions • Nonmetals tend to gain electrons so they become isoelectronic with its nearest noble gas neighbor located in the same period to the right O + 2e- O2- Oxygen atom Oxide ion 8e-, 6 valence e- 10e- [He]2s22p4 [He]2s22p6 or [Ne]
71. 71. 2.6 The Octet Rule Using the Octet Rule • The octet rule is very helpful in predicting the charges of ions in the representative elements • Transition metals still tend to lose electrons to become cations but predicting the charge is not as easy • Transition metals often form more than one stable ion – Iron forming Fe2+ and Fe3+ is a common example
72. 72. Examples Using the Octet Rule2.6 The Octet Rule • Give the charge of the • Which of the most probable ion following pairs of resulting from these atoms and ions are elements isoelectronic? – Ca – Cl-, Ar – Sr – Na+, Ne – S – Mg2+, Na+ – P – O2-, F-
73. 73. 2.7 Trends in the Periodic Table• Many atomic properties correlate with electronic structure and so also with their position in the periodic table – atomic size – ion size – ionization energy – electron affinity
74. 74. Atomic Size2.7 Trends in the Periodic • The size of an element increases, moving down from top to bottom of a group • The valence shell is higher in energy and Table farther from the nucleus traveling down the group • The size of an element decreases from left to right across a period • The increase in magnitude of positive charge in nucleus pulls the electrons closer to the nucleus
75. 75. 2.7 Trends in the Periodic Table Variation in Size of Atoms
76. 76. Cation Size2.7 Trends in the Periodic Cations are smaller than their parent atom • More protons than electrons creates an increased nuclear charge • Extra protons pull the remaining electrons closer to the nucleus Table • Ions with multiple positive charges are even smaller than the corresponding monopositive ions – Which would be smaller, Fe2+ or Fe3+? Fe3+ • When a cation is formed isoelectronic with a noble gas the valence shell is lost, decreasing the diameter of the ion relative to the parent atom
77. 77. Anion Size2.7 Trends in the Periodic Anions are larger than their parent atom. • Anions have more electrons than protons Table • Excess negative charge reduces the pull of the nucleus on each individual electron • Ions with multiple negative charges are even larger than the corresponding monopositive ions
78. 78. 2.7 Trends in the Periodic Relative Size of Select Ions and Their Parent Atoms Table
79. 79. 2.7 Trends in the Periodic Ionization Energy • Ionization energy - The energy required to remove an electron from an isolated atom • The magnitude of ionization energy Table correlates with the strength of the attractive force between the nucleus and the outermost electron • The lower the ionization energy, the easier it is to form a cation ionization energy + Na  Na+ + e-
80. 80. Ionization Energy of Select Elements2.7 Trends in the Periodic Table • Ionization decreases down a family as the outermost electrons are farther from the nucleus • Ionization increases across a period because the outermost electrons are more tightly held • Why would the noble gases be so unreactive?
81. 81. 2.7 Trends in the Periodic Electron Affinity • Electron affinity - The energy released when a single electron is added to an isolated atom Table • Electron affinity gives information about the ease of anion formation – Large electron affinity indicates an atom becomes more stable as it forms an anion Br + e–  Br– + energy
82. 82. 2.7 Trends in the Periodic Periodic Trends in Electron Affinity • Electron affinity generally Table decreases down a group • Electron affinity generally increases across a period