4 Chapter 6: The Periodic Table and Periodic Law Section 6.1 Development of the Modern Periodic TableAntoine Lavoisier( )Composed a list in 1790of the 23 elementsknown at the time.
5 The advent of electricity (used to break down compounds into their component elements- electrolysis) and the development of the spectrometer, used to identify elements lead to the discovery of many new elements (70 total by 1870)
26 Classifying the Elements MetalsAll elements to the left side (except H) of the stair-step line from B to At
27 Metals good conductor of heat and electricity shiny luster solid at room temperature (except Hg)Malleable and ductilehigh boiling and melting pointsHigh densities
28 Metals (continued) large atomic radius low ionization energy low electronegativity
29 Alkali Metals The group 1 elements (except for H) Most reactive group of metals !!!
30 Alkaline Earth Metals The group 2 elements Second most reactive group of metals
31 Transition ElementsThese are the group B elements (all are metals)
32 Transition MetalsTransition MetalsAlso known as the d-block elements
33 Inner transition Metals Located along the bottom of the periodic tableKnown as the lanthanide- actinide seriesLanthanides are also called rare earth elementsMany are radioactiveAlso know as the f- block elements
34 Non- Metals Occupy the upper right side of the periodic table (right of the stair-step line from B to At)
35 Non-Metals (Characteristics) Generally gasses at room temperaturePoor conductors of heat and electricityBrittle dull-looking when solidslower boiling and melting points than metals (except carbon)usually have lower densities than metals
36 Non-Metals ( Other Characteristics) high electronegativitySmall atomic radiushigher ionization energy than metals
37 Halogens Group 17 elements Most highly reactive group of the non-metalsDiatomic (Mr. BrINClHOF)
38 Nobel gases Group 18 elements Extreamly unreactive also called Inert gases
39 Metalloids Elements border the stair-step line from B to At Physical and chemical properties of both metals and non-metals
45 Section review 6.1 Laviosier,- Composed a list in 1790 of the 23 known elements.Newlands: Noticed that when elements were arranged by increasing atomic mass their properties repeated every eight element Law of Octaves.Mendeleev : Organized elements increasing atomic massElements with similar properties were arranged in columns- this was the first periodic tablepredicted the existence and properties of undiscovered elements and left spaces where they should goMoseley:By arranging elements by increasing atomic number, the problems with the order of the elements in the periodic table were solved.
51 Classification of the Elements By noting an atoms position on the periodic table, you can determine an its electron configuration and number of valence electrons
52 Valence Electrons electrons in an atoms highest principle energy level elements have similar chemical properties because they have the same number of valence electrons
53 Valence Electrons and Period The energy level of an elements valence electrons indicates the period on the periodic table in which it is foundEx. Li - found in period 2; valence electrons found in second energy levelEx. Gallium (Ga) found in period 4; valence electrons found in fourth energy level
54 Valence electrons and group number A main group (representative) element and the number of valence electrons it contains are related.Group 1 has one valence electron, Group 2 has 2 valence electrons and so on. ( Exception He is in group 18, but only has 2 valence electrons
55 Lewis electron dot structures illustrate this connection.
56 The s- p- d- and f- block elements The periodic table is divided into sections, or blocks representing the atoms energy sub levels being filled with valence electrons.
57 s- block elementsConsist of groups 1 and 2, H and He
58 s- block elements Valence electrons occupy only s orbitals group 1 elements have partially filled s orbitals s1group 2 elements have partially filled s orbitals s2
59 s- block elements Group 1 elements have partially filled s orbitals s1 Group 2 elements have filled s orbitals s2
60 p- block elementsAfter the s sublevel is full, the valence electrons next occupy the p sublevel with its three p orbitals.
61 p- block elementsThe p-block is comprised of groups 13 through 18, with filled or partially filled p orbitals
62 p- block elementsThere are no p block elements in the period 1, the first p-block element B is in period 2
63 p- block elementsTogether the s- and p- blocks compromise the main group elements.
64 p- block elementsGroup 18 are the Nobel gases that are unreactive with any other elements.
65 p- block elementsNobel gases have their highest PEL completely full, this is a very stable electron configuration.
66 d- block elementsLargest of the blocksContains the transition metals
67 d- block elementsCharacterized by a filled outermost s orbital and filled or partially filled d orbitalsd- block spans10 groups onthe periodic table
68 f- block elementsContains the inner transition metals (lanthanide-actinide series)
69 f- block elementsHave filled or partially filled outermost s orbital, and filled or partially filled 4f and 5f orbitals
70 F- blockElectrons do not fill their orbitals in a predictable manner.
71 F- block http://freezeray.com/flashFiles/discoveryDates.htm There are 7 seven f orbitals that hold a maximum of 14 electrons, so the f-block spans 14 columns of the periodic table.
72 Period 1: s-block elements only Period 1: s-block elements onlyPeriods 2 and 3: s- and p- block elementsPeriods 4 and 5: s-, p-, and d- block elementsPeriods 6 and 7: s-, p-, d- and f- block elements
73 Section 6.3 Periodic Trends As you move across a period or down a group elements tend to change in a predictable way, this is known as a periodic trend.
74 Atomic RadiusDefined as half the distance between adjacent nuclei in a crystal of the elementFor diatomic atoms AR is defined as half the distance between nuclei of identical atoms that are chemically bonded together.
75 Trends within periodsIn general there is a decrease in atomic radii as you move left-to-right across a period
76 Trends within periodsIn general there is a decrease in atomic radii as you move left-to-right across a period
77 The result is that the increased nuclear charge pulls the outermost electrons closer to the nucleus.
78 Trends within groupsAtomic radii generally increase as you move down a group.Electrons are added to higher PELs, so the outermost orbitals increase the size of the atom.The valence electrons are also farther from the nucleus
79 Largest: Na Smallest: S Largest: Xe Smallest: He No, with only this information, you will be unable to determine the specific groups and periods That the elements are in; so you can’t apply the periodic in atomic size to determine which element has the larger atomic radius.
80 No, with only this information, you will be unable to determine the soecific groups and periods That the elements are in; so you can’t apply the periodic in atomic size to determine which element has the larger atomic radius.
81 Ionic RadiusWhen atoms lose electrons they form positively charged ions, they always become smaller.
82 Ionic RadiusThis is due to the loss of electrons that are always valence electrons, thus the outer orbital is empty resulting in a smaller radius.
83 Ionic RadiusAtoms can gain or lose one or more electrons to form ions and gain a net charge.
84 Ionic RadiusAtoms can gain or lose one or more electrons to form ions and gain a net charge. An ion is an atom or bonded group of atoms that has a positive or negative charge.
85 Ionic RadiusWhen atoms gain electrons, they have a negative charge resulting in a larger radius.
86 Trends within periods Ionic Radius As you move left to right across the period the ionic radius decreases, then when you get to group 5A or 6A, the size of the larger negative ions also decreases.
87 Trends within groupsAs you move down a group, an ion’s outer electrons are in a higher PEL resulting in a gradual increase in size.
88 Ionic radiusAtoms can gain or lose one or more electrons to form ions and gain a net charge. An ion is an atom or bonded group of atoms that has a positive or negative charge.
89 Ionic radiusWhen atoms lose electrons they form positively charged ions, they always become smaller. This is due to the loss of electrons that are always valence electrons, Thus the outer orbital is empty resulting in a smaller radius.
90 Ionic radius When atoms gain electrons, they have a negative charge resulting in a smaller radius.
91 Ionic radius Trends within periods As you move left to right across the period the ionic radius decreases, then when you get to group 5A or 6A, the size of the larger negative ions also decreases.
92 Ionic radius Trends within groups As you move down a group, an ion’s outer electrons are in a higher PEL resulting in a gradual increase in size.
94 Ionization energyTo form a positive ion, an electron must be removed from a neutral atom.Ionization energy is defined as the energy required to remove an electron (from the gaseous state).The energy required to remove the first electron is called the first ionization energy (energy required to remove the second electron, is called the second ionization energy)
95 Ionization energyA high ionization energy value indicates that an atom has a strong hold on electrons.IE is a measure of how strongly an atom’s nucleus will hold onto valence electrons.Atoms with large IE are less likely to form positive ions than atoms with low IE.
96 Ionization energy Trends within periods First ionization energies generally increase as you move l to r across a period.
97 Ionization energy Trends within groups First ionization energies generally decrease as you move down a group.With the valence electrons farther away from the nucleus, less energy is required to remove them.
98 Octet RuleAtoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electron, (noble gas configuration).
99 Octet RuleThis is a very stable electron configuration. Elements to the right of the periodic table tend to gain electrons (form negative ions)Elements to the left of the periodic table tend to lose electrons (form positive ions to attain this configuration.
100 ElectronegativityThe relative ability of an atom to attract elements in a chemical bond.Expressed in terms of a numerical value of 3.98 or less.Units called Paulings (see Linus Pauling )
101 ElectronegativityThe relative ability of an atom to attract elements in a chemical bond.
102 Electronegativity Noble gases have been left out F is the most electronegative element (3.98); Cs and Fr are the least electronegative (.79 and .7)In a chemical bond the atoms with the greater electronegativity more strongly attract the bond’s electrons
103 Electronegativity Trends within groups and periods Electronegativity generally decreases as you move down a group and increases as you move left to right across a period.
104 Electronegativity Trends within groups and periods Elements with the higher electronegativity are at the upper right and elements with the lower electronegativity are at the bottom left of the periodic table.
107 AR increase down a group as the electrons are added to higher energy levels and inner core electrons shields the valence electrons from the increased nuclear charge. AR decrease across a period as increased nuclear charge coupled with unchanging shielding by inner core electrons pulls the valence electrons (being added to the same energy level) closer to the nucleus.Largest: antimony (Sb) Smallest: (nitrogen (N)
108 a. F c. Brb. Br d. FLithium’s second removed electron is an inner core electron, not a valence electron. Carbon’s fourth removed electron is still a valence electron.