2 Origins of the Periodic Table By the year 1700, only 13 elements had been identifiedScientific discovery led to a higher rate of element discoveryA logical organization of elements was needed for all the new elements
3 Early Organization J.W. Dobereiner (1829) organized elements in triads Triad – three elements with similar properties (ex: Cl, Br, I)J.R. Newlands (1864) organized elements in octavesOctave – repeating group of 8 elements
4 MendeleevDmitri Mendeleev (1869) arranged elements according to their propertiesMendeleev noticed that when the elements were arranged in order of increasing atomic mass, there was a repeating pattern to their propertiesThis is known as PeriodicityMendeleev left some spaces on his table blank, but was able to predict the properties of the unknown elements
6 Moseley Mendeleev’s table was imperfect – Te and I had to be reversed Henry Moseley (1913) arranged elements according to atomic numberThe periodic repetition of chemical and physical properties when elements are arranged by atomic number is known as the Periodic Law
7 Modern Periodic TableThe modern periodic table consists of Rows and ColumnsRows -HorizontalAlso known as PeriodsNumbered 1-7Columns -VerticalAlso known as Groups and FamiliesNumbered 1-18
8 Classifying ElementsThe elements on the periodic table can be simply classified by groupsGroups 1,2,13-18 (1A-8A) are known as the Representative Elements
9 Classifying ElementsGroups of representative elements have the same valence electrons and Oxidation StateOxidation State is how many electrons are gained or lost by an atom in a chemical reactionLost Electrons = Positive Oxidation StateGained Electrons = Negative Oxidation StateThink of Oxidation State as the charge of the ion
10 Driving ForceFull Energy Levels require lots of energy to remove their electrons.Noble Gases have full orbitals.Atoms behave in ways to achieve noble gas configuration
11 Classifying ElementsGroups 3-12 (3B-2B) , as well as the lanthanide and actinide series are known as Transition Metals
12 Metals The most common class of elements is Metals Metals become cationsWhat is a cation? How are they formed?Positively charged atom/positive oxidation state - Lose electronsMetals are generally solid (except Hg), conductive of heat and electricity, malleable, ductile, and shiny
13 Alkali Metals Group 1 elements are known as Alkali Metals Alkali metals include Li, Na, K, Rb, Cs, FrAlkali metals are generally dull, soft, and reactive – rarely found as free elements
14 Alkali Metals How many valence electrons do all Alkali Metals have? Write the noble gas configuration for each Alkali Metal[He]2s1[Ne]3s1[Ar]4s1[Kr]5s1[Xe]6s1[Rn]7s1How many valence electrons do all Alkali Metals have?OneWhat is the oxidation state of Alkali Metals?+1
15 Alkaline Earth MetalsGroup 2 elements are known as Alkaline Earth MetalsAlkaline earth metals include Be, Mg, Ca, Sr, Ba, and RaAlkaline earth metals are harder, denser, and stronger than alkali metalsLess reactive than alkali metals, but still rarely found as free elements
16 Alkaline Earth MetalsWrite the noble gas configuration for each Alkaline Earth Metal[He]2s2[Ne]3s2[Ar]4s2[Kr]5s2[Xe]6s2[Rn]7s2How many valence electrons do all Alkaline Earth Metals have?TwoWhat is the oxidation state of Alkaline Earth Metals?+2
17 Transition MetalsElements in groups 3-12 (3B-2B) are known as Transition MetalsTransition metals include Mn, Fe, Ag, Au, Mo, etc.Transition metals fill in the d orbital and often have multiple oxidation statesLanthanide and Actinide Series elements fill in the f orbitals – known as inner transition elements
18 MetalloidsElements that border the staircase on the periodic table are known as MetalloidsMetalloids include: B, Si, Ge, As, Sb, Te, Po, AtMetalloids have properties of both metals and nonmetals
19 NonmetalsNonmetals are found to the right of the staircase on the periodic tableNonmetals generally become anionsWhat is an Anion? How are they formed?Negatively charged atom/oxidation state - Gain electronsNonmetals are often gases or dull, brittle solidsNonmetals generally show poor conductivity, ductility, and malleability
20 Halogens Group 17 elements are known as Halogens Halogens include F, Cl, Br, and IHalogens are the most reactive nonmetals – often found in compounds
21 Halogens Write the noble gas configuration for each Halogen [He]2s22p5 [Ne] 3s23p5[Ar] 4s23d104p5[Kr] 5s24d105p5How many valence electrons do all Halogens have? Oxidation State?Seven / -1Why are the Halogens the most reactive non-metals?They are 1 electron short of having an octet.
22 Noble Gases Elements in group 18 are known as Noble Gases Noble Gases include He, Ne, Ar, Kr, Xe, RnNoble gases are extremely unreactive
23 Noble Gases How many valence electrons do all Noble Gases have? Eight Write the electron configuration for each Noble Gas1s2[He]2s22p6[Ne]3s23p6[Ar]4s23d104p6[Kr]5s24d105p6[Xe]6s25d106p6How many valence electrons do all Noble Gases have?EightWhy are Noble Gases so unreactive?They contain a full octet – atoms gain/lose electrons to achieve noble gas notation
24 Other GroupsAll other groups can be identified by the top most element in that group.Ex: Group 15 can be called the Nitrogen GroupOxidation State: -3Q: What is another name for Group 16?A: Oxygen groupQ: Oxidation StateA: -2
25 Periodic TrendsThe elements on the periodic table show repeating trends related to electron configuration
27 Atomic RadiusThe Atomic Radius is ½ the distance between nuclei of bonded atoms from the same elementAtomic radius decreases from left to right across a periodAtomic radius increases from top to bottom in a period
28 Why?Not changing energy level, but increasing nuclear force (more positive charge in nucleus)
29 Ionization EnergyIf an atom is becoming an ion, it is gaining or losing electrons in an effort to have an octet (8 valence electrons)
30 Ionization EnergyThe energy required to remove an electron from an atom is called Ionization Energy
31 Ionization Energy1st Ionization Energy- energy required to remove 1st electron from an atom2nd Ionization Energy- energy required to remove 2nd electron from an atom2nd Ionization Energy is ALWAYS higher than the 1st3rd Ionization Energy- energy required to remove 3rd electron from an atom3rd Ionization Energy is ALWAYS higher than the 1st or 2nd
32 Ionization Energy IE Decreases as you move down a group Why? Electron is further away
33 Ionization Energy IE Increases as you move across a period Why? You are in the same energy level but have more nuclear charge
34 Ionization EnergyFull Energy Levels require lots of energy to remove their electrons.Noble Gases have full orbitals.Atoms behave in ways to achieve noble gas configuration.
35 Ionization Energy Write the electron configuration for Be 1s22s2 How many valence electrons does Be have?2Why is the ionization energy low?It is easier for Be to lose those 2 valence electrons than it is to gain 6. Therefore, it has a low ionization energy.
36 Ionization EnergyMove across the period. Write the electron configuration for F.1s22s22p5How many valence electrons does F have?7Why is the ionization energy high?It is easier for F to gain 1 valence electron than is it for it to lose 7. Therefore, its’ ionization energy (energy to lose an electron) is high
38 Electron AffinityElectron affinity is the energy change associated with adding an electron to a gaseous atom.Easiest to add to group 7A (halogens).Why?Gets them to full octet.Increase from left to right: atoms become smaller, with greater nuclear charge.Decrease as we go down a group.
39 Ionic SizeCations are smaller than the atoms from which they form (less electrons)Anions are larger than the atoms from which they form (more electrons)
40 Ionic SizeAcross the period, nuclear charge increases so they get smaller.Energy level changes between anions and cations.N3-O2-F1-B3+Li1+C4+Be2+
41 ElectronegativityElectronegativity is the ability for an atom to attract electrons in a compoundElectronegativity increases from left to right in a periodElectronegativity decreases from top to bottom in a group
42 ElectronegativityWe do not consider noble gases when talking about electronegativity because they do not bond.What is the most electronegative element?Fluorine
43 Electronegativity Write the electron configuration for Li 1s22s1 How many valence electrons does Li have?1Why is the electronegativity low??It is easier for Li to lose 1 valence electrons than it is to gain 7. It has a low electronegativity because it would be difficult for Li to attract 7 electrons
44 ElectronegativityMove across the period. Write the electron configuration for O.1s22s22p4How many valence electrons does O have?6Why is the electronegativity high?It is easier for O to gain 2 valence electrons than is it for it to lose 6. Electronegativity is high because it can gain electrons more easily than it can lose them.