Chapter 6 The Periodic Table Why are the elements cobalt phosphorus and sulfur essential for maintaining law and order? Because when they are put together, they’re CoPS

6.1 Organizing the Elements u OBJECTIVES: Explain how elements are organized in a periodic table Compare early and modern periodic tables Identify three broad classes of elements

Section 6.1 Organizing the Elements u A few elements, such as gold and copper, known for thousands of years u Yet only about 13 had been identified by the year u As more elements discovered, chemists realized a way to organize the elements was needed.

6.1 Organizing the Elements u Chemists used properties of elements to sort them into groups. u In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties One element in each triad had properties intermediate of other two Worked well for many elements, but not all

Mendeleev’s Periodic Table u By mid-1800s, about 70 elements known u Dmitri Mendeleev, a Russian chemist and teacher, arranged elements in order of increasing atomic mass u Mendeleev noticed that if he grouped elements a certain way, certain chemical characteristics appeared periodically on the table u Thus, the Periodic Table

Mendeleev u Mendeleev left “blanks” on his table where he thought undiscovered elements should fit When some of these elements were discovered, many of his predictions proved to be spot on u But, there were problems: Co and Ni; Ar and K; Te and I He’d switched them on his table because he didn’t know about isotopes

A better arrangement u Mendeleev’s mistake: arranging elements based on atomic mass instead of z so he’d switched Co and Ni because the atomic mass of Co is slightly higher than Ni, due to Co having heavier isotopes (more neutrons) u In 1913, Henry Moseley, a British physicist, arranged elements according to increasing atomic number u The symbol, atomic number & mass are basic items included on most tables

Spiral Periodic Table

The Periodic Law says: u When elements are arranged in order of increasing atomic number, a periodic repetition of physical and chemical properties occurs u Horizontal rows = periods There are 7 periods u Vertical columns = groups (or families) Similar physical & chemical props –Especially main block elements (groups 1A - 8A)

Classes of Elements 1. Metals: electrical conductors, have luster, ductile, malleable 2. Nonmetals: brittle non-lustrous solids, or gases, one liquid (Br), all poor conductors of heat & electricity 3. Metalloids: border zig-zag line on both sides and have properties intermediate between metals and nonmetals

Checking Understanding u What did Mendeleev base his table of elements upon? u What caused Mendeleev to mistakenly switch Co and Ni; Ar and K; Te and I on his original table? u Complete the following table (Note charges on ions): ElementMetal?ZMass #neutronselectrons 72 GeMetalloid Ca O 127 I − 59 Co Ar

6.2 Classifying the Elements u OBJECTIVES: Describe the information provided in a periodic table Classify elements based on electron configurations Distinguish representative elements (main block) and transition metals

Squares in a Periodic Table The periodic table displays symbols and names of elements, along with information about the structure of their atoms: Atomic number and atomic mass Symbols usually color coded solid; gas; liquid

Based on the average atomic mass and the atomic number, what is likely the most common isotope of Na?

Groups of Elements - Families u Group 1A – alkali metals Forms a “base” (or alkali) when reacting (very violently!) with water Na + 2 H 2 O  Na OH − + H 2 u Group 2A – alkaline earth metals Also form bases with water do not dissolve well and melt at very high T, hence usually solid (“earth”) u Group 7A – halo-gens (= “salt-forming”)

Electron Configurations in Groups u Elements can be sorted into 4 different groupings based on their electron configurations: 1) Noble gases 2) Representative elements 3) Transition metals 4) Inner transition metals Let’s now take a closer look at these.

Electron Configurations in Groups 1. Main Block Elements = Groups 1A - 8A a) Representative Elements Groups 1A - 7A i. Display wide range of properties, thus a good “representative” sample of all matter ii. There are metals, nonmetals, and metalloids; Many are solid, others are gases or liquids iii. Their outermost electrons are in s and p subshells that are NOT filled b) Noble Gases: so called because they rarely take part in a reaction; very stable (exist as free atoms in nature) i. Noble gases have their outer s and p sublevels completely full (the most stable arrangement)

Electron Configurations in Groups 2) Transition (d-Block) metals in the “B” columns of the U.S. style periodic table Electron configuration has the outer s sublevel full (usually), and all or part of d subshell filled Block “transitions” between the metal and nonmetal area of table Examples include familiar metals: gold, copper, silver, iron

Electron Configurations in Groups 3) Inner Transition Metals located below main body of the table, in two rows Electron configuration has the outer s sublevel full, and f sublevel is partly to completely filled Formerly called “rare-earth” elements, but this is not true because some are quite abundant We won’t be concerned with these much in this class

1A 2A3A4A5A6A 7A 8A u Main Block in color u Groups 1A-7A are representative elements

The group B are called the transition elements u These are called the inner transition elements, and they belong here

Group 1A are the alkali metals (but NOT H) Group 2A are the alkaline earth metals H

u Group 8A are the noble gases u Group 7A is called the halogens

1s11s1 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6 s 2 4f 14 5d 10 6p 6 7s 1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Do you notice any similarity in the configurations of the alkali metals?

He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86 1s21s2 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 Do you notice any similarity in the configurations of the noble gases?

u Alkali metals all end in s 1 u Alkaline earth metals all end in s 2 really should include He, but it fits better in a different spot, since He has the properties of the noble gases, and has a full outer level of electrons. s2s2 s1s1 Elements in the s - blocks He

Transition Metals - d block d1d1 d2d2 d3d3 s1d5s1d5 d5d5 d6d6 d7d7 d8d8 s 1 d 10 d 10 Note exceptions to electron configuration rules

The P-block p 1 p 2 p 3 p 4 p 5 p 6

u Each row (or period) is the energy level for s and p orbitals (d block is one level lower in each row) Period Number 3d  4d 

u f orbitals start filling at 4f, and are 2 less than the period number f 5f

6.3 Periodic Trends u OBJECTIVES: Describe and explain trends among the elements for atomic size Explain how ions form Describe and explain periodic trends for first ionization energy, ionic size Describe the periodic trend for electronegativity

Trends in Atomic Size u How do we measure? electron clouds of atoms don’t have a definite edge u Chemists get around this by measuring more than 1 atom at a time

Atomic Size u Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule. } Radius

3 Factors Affect Periodic Trends 1. Energy Level  Electrons in higher energy levels are farther away from the nucleus  Each successive level builds on top of lower ones 2. Nuclear Charge (# protons)  More + charge pulls electrons in closer. (+ and – attract each other) 3. Shielding effect  “core” electrons near nucleus shield outer electrons from protons’ “pull”  The higher the energy level  greater shielding

#1. Atomic Size - Group trends u As we go down a group, each atom has another energy level (= more electrons, farther from nucleus) u so the atoms get bigger. H Li Na K Rb

#1. Atomic Size - Period Trends u Going from left to right across a period, the size gets smaller. u Electrons are in the same energy level. u But, there is more nuclear charge. u Outermost electrons are pulled closer. NaMgAlSiPSClAr

Atomic Number Atomic Radius (pm) H Li Ne Ar 10 Na K Kr Rb 3 Period 2

Ions u Some compounds are composed of particles called “ions” An ion is a atom (or group of atoms) that has a positive or negative charge u Atoms are neutral because the number of protons equals electrons Positive and negative ions are formed when electrons are transferred (lost or gained) between atoms

Ions u Metals tend to LOSE electrons, from their outer energy level Sodium loses one: there are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation” Charge is written as a number followed by a + or – sign: Na 1+ or Na + Now named a “sodium ion”

Ions u Nonmetals tend to GAIN one or more electrons Oxygen often gains two electrons Protons (8) no longer equals electrons (10), so a charge of -2 O 2– is re-named a “oxide ion” Negative ions are called “anions”

#2. Trends in Ionization Energy u Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom) u Removing one electron makes a 1+ ion u The energy required to remove only the first electron is called the first ionization energy u Energy needed to remove a second electron = second ionization energy u Remove a third = third ionization energy u And so on…

Ionization Energy u The second ionization energy is always greater than first IE There are more protons than electrons after the first is removed, so the nucleus pulls the rest closer u The third IE (energy to remove a third electron) is even greater Removing an electron from a 2+ ion

SymbolFirstSecond Third H He Li Be B C N O F Ne Table 6.1, p. 173

SymbolFirstSecond Third H He Li Be B C N O F Ne Why did these values increase so much ?

What factors determine IE u Greater nuclear charge = greater IE More protons, more difficult to remove e - u Greater distance from nucleus decreases IE Farther from “pull” of protons + shielding effect of e - in lower energy levels u Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE.

Shielding u Nucleus has weak attraction to electron in outermost energy level due to repulsive forces of all the other energy levels in between. u Removing a second e – from the same period requires only a little more energy because shielding is the same

SymbolFirstSecond Third Ne Na Mg Al Why did these values increase so much ? The big jumps in IE seen above occur because that second electron removed from Na and the third removed from Mg are being taken from a lower energy level (a core electron rather than a valence electron)

Ionization Energy - Group trends u As you go down a group, the first IE decreases because... The electron is further away from the attraction of the nucleus, and There is more shielding.

Ionization Energy - Period trends u All the atoms in the same period have the same energy level = same shielding BUT, increasing nuclear charge u So IE generally increases from left to right u Exceptions occur (of course), but can be explained in light of electron configurations

Trends in Ionic Size: Cations u Cations form by losing electrons u Cations are smaller than the atom they came from – not only do they lose electrons, they often lose an entire energy level u Metals form cations u Cations of representative elements have the noble gas configuration before them

Ionic size: Anions u Anions form by gaining electrons. u Anions are bigger than the atom they came from – have the same energy level, but more electrons to attract with the same nuclear charge u Nonmetals form anions. u Anions of representative elements generally have a noble gas configuration after they ionize.

Configuration of Ions u Main block ions have noble gas configurations ( = a full outer level) u Na atom is: 1s 2 2s 2 2p 6 3s 1 u Forms a 1+ sodium ion: 1s 2 2s 2 2p 6 u Same configuration as neon. u Metals form ions with the configuration of the noble gas before them - they lose electrons.

Configuration of Ions u Non-metals form ions by gaining electrons to achieve a noble gas configuration u The ion’s noble gas configuration is the same as the noble gas that comes after them in the periodic table F = 1s 2 2s 2 2p 5 F – = 1s 2 2s 2 2p 6

Ion Group trends u Each step down a group is adding an energy level u Ions get bigger as you go down, because of the additional energy level Same as with atoms Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

Ion Period Trends u Across the period from left to right, the nuclear charge increases - so they get smaller. u Notice the energy level changes between anions and cations. Li 1+ Be 2+ B 3+ C 4+ N 3- O 2- F 1-

Size of Isoelectronic ions u Iso- means “the same” u Isoelectronic ions have the same # of electrons u Al 3+ Mg 2+ Na 1+ Ne F 1- O 2- and N 3- all have 10 electrons u all have the same configuration: 1s 2 2s 2 2p 6 (which is = noble gas neon)

Size of Isoelectronic ions? u Positive ions that have more protons would be smaller (more protons would pull the same # of electrons in closer) Al 3+ Mg 2+ Na 1+ Ne F 1- O 2- N

#3. Trends in Electronegativity u Electronegativity is the tendency for an atom to attract shared electrons when it is chemically combined with another element. u They share electrons, but how equally do they share them? u An element with a high electronegativity means it pulls the electron towards itself strongly!

Electronegativity Group Trend u The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has u Valence electrons farther from the nucleus, so less tightly held u Thus, more “willing to share” = Low electronegativity

Electronegativity Period Trend u Metals let their electrons go easily Thus, low electronegativity u As you move from left to right metallic character decreases, until you get to the non-metals at the far right of the table u Adding electrons makes non-metals more stable, so they often take them away from other atoms = High electronegativity

The arrows indicate the trend: Ionization energy and Electronegativity INCREASE in these directions

Atomic size increases in these directions:

First Ionization energy Atomic number He u He has a greater IE than H. u Both elements have the same shielding since electrons are only in the first level u But He has a greater nuclear charge H

First Ionization energy Atomic number H He Li has lower IE than H more shielding further away l These outweigh the greater nuclear charge Li

First Ionization energy Atomic number H He Be has higher IE than Li same shielding l greater nuclear charge Li Be

First Ionization energy Atomic number H He B has lower IE than Be similar shielding greater nuclear charge l But 2p e – removed from B is higher energy than the 2s removed from Be l This makes up for the greater nuclear charge of B, and thus the IE is slightly lower than Be Li Be B

First Ionization energy Atomic number H He Li Be B C

First Ionization energy Atomic number H He Li Be B C N

First Ionization energy Atomic number H He Li Be B C N O u Pattern broken again, because electron removed from O comes out of a filled p orbital u The added repulsion of another electron in the same orbital makes it easier to remove one

First Ionization energy Atomic number O F H He Li Be B C N

First Ionization energy Atomic number H He Li Be B C N O F Ne u Ne has a lower IE than He u Both are full, u Ne has more shielding u Greater distance

First Ionization energy Atomic number H He Li Be B C N O F Ne Na has a lower IE than Li Both are s 1 Na has more shielding l Greater distance Na

First Ionization energy Atomic number