3 Until 1800’s – no clear system The Search for a Periodic TableUntil 1800’s – no clear systemElements grouped by similar properties or atomic massIn 1829, J.W. Döbereiner classified some elements into groups of three, which he called triads.1600’s – chemistry still tied to alchemyBelief that urine could be changed to gold led to discovery of matches1700’s – elan vital – force of lifeThen mid 1700’s = Lavoisier1800’s – no clear system used individual abbreviations, symbols. Formulas meant different things to different people1850’s – Mendeleev – grouped by atomic mass or common properties (solids, gases)
4 Döbereiner’s TriadsThe elements in a triad had similar chemical properties, and their physical properties varied in an orderly way according to their atomic masses.ElementAtomic mass (g)Density (g/mL)Melting point (C)Boiling point (C)Chlorine35.5-101-34Bromine79.93.12-759Iodine1274.93114185
5 Density increases with increasing atomic mass. Döbereiner’s TriadsElementAtomic mass (g)Density (g/mL)Melting point (C)Boiling point (C)Chlorine35.5-101-34Bromine79.93.12-759Iodine1274.93114185Density increases with increasing atomic mass.The concept of triads suggested that the properties of an element are related to its atomic mass.
6 Which of the Dobereiner triads shown are still listed in the same column of the modern periodic table?Triad 1Triad 2Triad 3LiMnSNaCrSeKFeTeTriad 1 and triad 3
7 organized the elements according to increasing atomic mass. Mendeleev’s Periodic TableThe Russian chemist, Dmitri Mendeleev, developed a periodic table of elements.organized the elements according to increasing atomic mass.
8 Mendeleev’s Periodic Table Mendeleev later developed an improved version of his table with the elements arranged in horizontal rows.
9 Mendeleev’s Periodic Table Patterns of changing properties repeated for the elements across the horizontal rows.Elements in vertical columns have similar properties.
10 Periodicity is the tendency to recur at regular intervals. Mendeleev’s Periodic Tableproperties of the elements repeat in an orderly way from row to row of the table.This repeated pattern is an example of periodicity in the properties of elements.Periodicity is the tendency to recur at regular intervals.
11 Mendeleev’s Periodic Table In order to group elements with similar properties in the same columns, Mendeleev had to leave some blank spaces in his table.He suggested that these spaces represented undiscovered elements.***Mendeleev correctly predicted the properties of several undiscovered elements.Why is this important?
15 The Modern Periodic Table the basis for ordering the elements in the table is the atomic number, not atomic mass.The atomic number of an element is equal to the number of protons in the nucleus.Each row (except the first) begins with a metal and ends with a noble gas.
16 The Modern Periodic Table In between, the properties of the elements change in an orderly progression from left to right.This regular cycle illustrates periodicity in the properties of the elements.
17 The Modern Periodic Table periodic law - physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number
18 Use the periodic table to separate these 12 elements into 6 pairs fo elements having similair properties. Ca, K, Ga, P, Si, Rb, B, Sr, Sn, Cl, Bi, BrCaKGaPSiClSrRbBBiSnBr
19 Layout of the Periodic Table Have them get out their periodic tables for labeling
20 Layout of the periodic table A group, also called a family, consists of the elements in a vertical column.
21 Groups are numbered 1 – OR IA – VIIIA for main group elements and IB – VIIIB for transition elementsMake sure to point out the weird numbering with the transitions elements
22 As you move left to right across a period the number of valence electrons increases by one
23 Elements in the same group have same number of valence electrons and similar properties
24 A period consists of the elements in a horizontal row
25 Periods are numbered 1-7 and each new row begins a new energy level
26 The elements in the middle are called transition elements
31 the halogens in Group 17 (VIIA) -from the Greek words for “salt former” , compounds that halogens form with metals are salt-like.
32 the noble gases in Group 18 (VIIIA) – full outer shell (8 valence electrons), generally unreactive
33 In the periodic table, two series of elements are placed below the main body of the table. The elements in these two series are known as the inner transition elements.
34 The first series of inner transition elements is called the lanthanides because they follow element number 57, lanthanum.Because of their natural abundance on Earth is less than 0.01 percent, the lanthanides are sometimes called the rare earth elements.
35 The second series of inner transition elements are the actinides All of the actinides are radioactive, and all beyond uranium (92) are man made (synthetic).
38 Elements are classified as metals, metalloids, or nonmetals on the basis of their physical and chemical properties.The majority of the elements are metals (solids). They occupy the entire left side and center of the periodic table.
39 Physical States and Classes of the Elements Nonmetals occupy the upper-right-hand corner. – green, yellow, orange
40 Physical States and Classes of the Elements Metalloids are located along the staircase boundary between metals and nonmetals. - purple
41 MetalsMetals are elements that have luster, conduct heat and electricity, and usually bend without breaking.All metals except mercury are solids at room temperature; in fact, most have extremely high melting points.Click box to view movie clip.
42 MetalsWith the exception of tin, lead, and bismuth, metals have one, two, or three valence electrons.The periodic table shows that most of the metals (coded blue) are not main group elements.
43 NonmetalsMost nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are brittle when solid.Their melting points tend to be lower than those of metals.Many are gases at room temperature
44 With the exception of carbon, nonmetals have five, six, seven, or eight valence electrons.
45 Properties of Metals and Nonmetals Copy the chart
46 MetalloidsMetalloids have some chemical and physical properties of metals and other properties of nonmetals. - purpleIn the periodic table, the metalloids lie along the border between metals and nonmetals.
47 some metalloids are semiconductors A semiconductor is an element that does not conduct electricity as well as a metal, but does conduct slightly better than a nonmetal.Some metalloids such as silicon, germanium (Ge), and arsenic (As) are semiconductors.Silicon’s semiconducting properties made the computer revolution possible.
49 Periodic Properties of the Elements The electron structure of an atom determines many of its chemical and physical properties.Understanding the relationship between electron configuration and position in the periodic table enables you to predict the properties of the elements and the outcome of many chemical reactions.
51 Atomic Sizesize of an atom INCREASES in any group as you go DOWN the column because the valence electrons are in energy levels farther from the nucleus.
52 “shielding effect” – electrons in energy levels closer to the nucleus “shield” the valance electrons from the positive pull of the nucleus
53 The shielding effectIncreases down a group because electrons are being added to higher energy levelsThere is no shielding effect as you go across a period because electrons are being added to the same principal energy level
54 size of an atom DECREASES in any period as you go to the RIGHT in any row because there is an increased nuclear (+) charge pulling e- in tighter.
55 Atomic Radius Why larger going down? Why smaller to the right? Higher energy levels have larger orbitalsShielding - core e- block the attraction between the nucleus and the valence e-Why smaller to the right?Increased nuclear charge without additional shielding pulls e- in tighter
57 ExamplesWhich atom has the larger radius?Be or BaCa or BrBaCa
58 For each of the following pairs, predict which atom is larger. a. Mg, SrSrb. Sr, SnSrc. Ge, SnSnd. Ge, BrGee. Cr, WW
59 Octet Rulereactivity of atoms is based on achieving a complete octet of valence electrons(8/8)Everybody wants to be like a noble gas!Ne
60 Atoms achieve noble gas configuration by gaining or losing their valence electrons An ion is an atom or group of atoms that has a charge because of the loss or gain of electrons.
61 cation - An ion that has LOST an e- and now has a positive (+) charge anion – an ion that has GAINED an e- and now has a negative (-) charge
62 Common Ion Charges aka oxidation number 1+1-3+3-2-2+
63 Group IVA 4 can lose or gain Group VA 5 gains 3 Group VIA 6 gains 2 GROUP #: VALENCE # WHEN FORMING IONS:OUT OF 8:Group IA loses 1Group IIA loses 2Group IIIA loses 3Group IVA 4 can lose or gainGroup VA gains 3Group VIA gains 2Group VIIA gains 1Group VIIIA 8 does not form ions
64 Ionic Sizepositive ions + (cations) acquire the configuration of the noble gas in the preceding period.the outermost electrons of the ion are in a lower energy level than the valence electrons of the neutral atom.
65 The electrons that are not lost by the atom experience a greater attraction to the nucleus and pull together in a tighter bundle with a smaller radius.all cations ions have smaller radii than their corresponding atoms.
66 anions acquire the electron configuration of the noble gas at the end of its period. But the nuclear charge doesn’t increase with the number of electrons.
67 In the case of fluorine, a nuclear charge of 9+ must hold ten electrons in the F– ion; all the electrons are held less tightlythe radius of the anion is larger than the neutral atom.
72 group trends(first) ionization energy decreases from top to bottom along a groupreason: outermost electron is farther and farther from the nucleus in larger atoms, so it is more easily removed
73 periodic trends(first) ionization energy increases from left to right in a periodreason: ―nuclear charge(+) increases; more attraction between electrons and protons
74 Successive Ionization Energies Large jump in I.E. occurs when a CORE e- is removed.Mg 1st I.E kJ2nd I.E. 1,445 kJCore e- 3rd I.E. 7,730 kJ
75 Successive Ionization Energies Large jump in I.E. occurs when a CORE e- is removed.Al 1st I.E kJ2nd I.E. 1,815 kJ3rd I.E. 2,740 kJCore e- 4th I.E. 11,600 kJ
76 ExamplesWhich atom has the higher 1st I.E.?N or BiBa or NeNNe
77 For each of the following pairs, predict which atom has the higher first ionization energy. a. Mg, NaMgb. S, OOc. Ca, BaCad. Cl, ICle. Na, AlAlf. Se, BrBr
78 Periodic Trends in Electronegativity electronegativity— tendency of an atom to attract electrons.noble gases do not have electronegativity valueschemical bonds are determined by electronegativity differences between the bonding partners
79 electronegativity trends are not completely regular fluorine = most electronegative element with a value of 4.0 (smallest anion formed)cesium = least electronegative element (largest cation formed)
80 electronegativity decreases from top to bottom in a group
81 electronegativity increases from left to right in a period
88 Electrons are arranged in orbits around the nucleus Electrons in AtomsNiels Bohr.Electrons are arranged in orbits around the nucleusThe energy level of an electron is the region around thenucleus where theelectron is likely tobe moving.
89 modern 3-D electron-cloud model - probability model Heisenberg Uncertainty Principle — it is not possible to know both the exact position and velocity of an object simultaneously
90 Modern electron cloud model orbitals are areas of high probability (~95%) of finding electrons
91 Electrons can change energy level, by absorbing energy Electrons can change energy level, by absorbing energy. When an electron absorbs a quantum of energy, it moves up to a higher energy level.When the electron falls from a higher energy level to a lower energy level, energy is released, and we see light
92 Energy levels have sublevels —divisions within an energy level 1) many similar energy states grouped together in a level2) different shapes: spherical, dumbbell, cloverleaf
93 There are 4 sublevelss, p, d, f(s p d f stand for sharp, principal, diffuse, fundamental)maximum number of e- in a principalenergy level = 2n 2n = principal quantum number= electron energy level or ― “shell” (period) numbern = 1, 2, 3, 4, 5, 6, 7
94 electron maximums in the sublevels s can hold 2 e-p can hold 6 e-d can hold 10 e-f can hold 14 e-
95 Electrons fill orbitals in a certain way electron configuration - a specific electron arrangement in orbitals
96 Electron configuration: General Rules Pauli Exclusion PrincipleEach orbital can hold 2 electrons with opposite spins.
97 Aufbau PrincipleElectrons fill the lowest energy orbitals first.“Lazy Tenant Rule”
98 WRONG RIGHT Hund’s Rule Within a sublevel, place one e- per orbital before pairing them.“Empty Bus Seat Rule”WRONGRIGHT
99 Different sections of the periodic table correspond to the different sublevels Groups IA & IIA = s blockGroups IIIA – VIIIA = p blockTransition = d blockInner transition = f block
101 Diagonal rule -to help us remember the order in which energy level subshells fill - follow the arrows1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p
102 1 s2 2s2 2p4 O Example 1s2 2s2 2p4 8e- Orbital Diagram Electron Configuration =1 s2 2s2 2p4
103 1s2 2s2 2p4the sum of the superscripts = the atomic number of the elementsuperscripts are NOT exponents (nothing is being squared, etc.)
104 *** valence configurations will be s OR s and p ***
105 Longhand Configuration 1s2 2s2 2p6 3s2 3p4 S 16e- Condensed (Abbreviated) Electron Configurationsuse the previous Noble Gas as the starting point in brackets, then finish the configurationLonghand Configuration1s22s22p63s23p4S 16e-Core ElectronsValence ElectronsShorthand ConfigurationS 16e- [Ne] 3s2 3p4