Periodicity

Classification of the Elements u OBJECTIVES: Explain why you can infer the properties of an element based on those of other elements in the periodic table.

Classification of the Elements u OBJECTIVES: Use electron configurations to classify elements as noble gases, main group elements, transition metals, or inner transition metals.

Periodic Table Revisited u Russian scientist Dmitri Mendeleev taught chemistry in terms of properties. u Mid 1800’s - molar masses of elements were known. u Wrote down the elements in order of increasing mass. u Found a pattern of repeating properties.

Mendeleev’s Table u Grouped elements in columns by similar properties in order of increasing atomic mass. u Found some inconsistencies - felt that the properties were more important than the mass, so switched order. u Also found some gaps. u Must be undiscovered elements. u Predicted their properties before they were found.

The modern table u Elements are still grouped by properties. u Similar properties are in the same column. u Order is by increasing atomic number. u Added a column of elements Mendeleev didn’t know about. u The noble gases weren’t found because they didn’t react with anything.

u Horizontal rows are called periods u There are 7 periods

Vertical columns called groups Elements are placed in columns by similar properties Also called families

1A 2A3A4A5A6A 7A 8A 0 u The elements in the A groups are called the representative elements outer s or p filling

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

u Group 1A are the alkali metals u Group 2A are the alkaline earth metals

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

Why? u The part of the atom another atom sees is the electron cloud. u More importantly the outside orbitals. u The orbitals fill up in a regular pattern. u The outside orbital electron configuration repeats. u The properties of atoms repeat.

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

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

u Alkali metals all end in s 1 u Alkaline earth metals all end in s 2 u really should include He, but it fits better later. u He has the properties of the noble gases. s2s2 s1s1 S- block

Transition Metals -d block d1d1 d2d2 d3d3 s1d5s1d5 d5d5 d6d6 d7d7 d8d8 s 1 d 10 d 10

The P-block p1p1 p2p2 p3p3 p4p4 p5p5 p6p6

F - block u inner transition elements

u Each row (or period) is the energy level for s and p orbitals

u d orbitals fill up after previous energy level, so first d is 3d even though it’s in row d

u f orbitals start filling at 4f f 5f

Writing electron configurations the easy way

Electron Configurations repeat u The shape of the periodic table is a representation of this repetition. u When we get to the end of the column the outermost energy level is full. u This is the basis for our shorthand.

The Shorthand u Write symbol of the noble gas before the element, in [ ]. u Then, the rest of the electrons. u Aluminum’s full configuration: 1s 2 2s 2 2p 6 3s 2 3p 1 u previous noble gas Ne is: 1s 2 2s 2 2p 6 u so, Al is: [Ne] 3s 2 3p 1

More examples u Ge = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 Thus, Ge = [Ar] 4s 2 3d 10 4p 2 u Hf = 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 2 Thus, Hf = [Xe]6s 2 4f 14 5d 2

The Shorthand Again Sn- 50 electrons The noble gas before it is Kr [ Kr ] Takes care of 36 Next 5s 2 5s 2 Then 4d 10 4d 10 Finally 5p 2 5p 2

Periodic Trends u OBJECTIVES: Interpret group trends in atomic radii, ionic radii, ionization energies, m.p., b.p., electronegativity and chemical properties

Trends in Atomic Size u First problem: Where do you start measuring from? u The electron cloud doesn’t have a definite edge. u They get around this by measuring more than 1 atom at a time.

Atomic Size u Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

Trends in Atomic Size u Influenced by three factors: 1. Energy Level Higher energy level is further away. 2. Charge on nucleus More charge pulls electrons in closer. u 3. Shielding effect e e repulsion

Group trends u As we go down a group... u each atom has another energy level, u so the atoms get bigger. H Li Na K Rb

Periodic Trends u As you go across a period, the radius gets smaller. u Electrons are in same energy level. u More nuclear charge. u Outermost electrons are closer. NaMgAlSiPSClAr

Overall Atomic Number Atomic Radius (nm) H Li Ne Ar 10 Na K Kr Rb

Trends in Ionization Energy u The amount of energy required to completely remove a mole of electrons from a mole of gaseous atoms. u Removing an electron makes a +1 ion. u The energy required to remove (1 mole of) the first electron is called the first ionization energy.

Ionization Energy u The second ionization energy is the energy required to remove (1 mole of) the second electron(s). u Always greater than first IE. u The third IE is the energy required to remove a third electron. u Greater than 1st or 2nd IE.

SymbolFirstSecond Third H He Li Be B C N O F Ne

SymbolFirstSecond Third H He Li Be B C N O F Ne

What determines IE u The greater the nuclear charge, the greater IE. u Greater distance from nucleus decreases IE u Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. u Shielding effect

Shielding u The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. u Second electron has same shielding, if it is in the same period

Group trends u As you go down a group, first IE decreases because... u The electron is further away. u More shielding.

Periodic trends u All the atoms in the same period have the same energy level. u Same shielding. u But, increasing nuclear charge u So IE generally increases from left to right. u Exceptions at full and 1/2 full orbitals.

First Ionization energy Atomic number He u He has a greater IE than H. u same shielding u greater nuclear charge H

First Ionization energy Atomic number H He Li has lower IE than H more shielding further away l outweighs 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 same shielding greater nuclear charge l p orbital is slightly more diffuse and its electron easier to remove 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 Breaks the pattern, because removing an electron leaves 1/2 filled p orbital

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

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

Driving Force u Full Energy Levels require lots of energy to remove their electrons. u Noble Gases have full orbitals. u Atoms behave in ways to achieve noble gas configuration.

2nd Ionization Energy u For elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected. u True for s 2 u Alkaline earth metals form 2+ ions.

3rd IE u Using the same logic s 2 p 1 atoms have an low 3rd IE. u Atoms in the aluminum family form 3+ ions. u 2nd IE and 3rd IE are always higher than 1st IE!!!

Trends in Electron Affinity u The energy change associated with adding an electron to a gaseous atom. u Easiest to add to group 7A. u Gets them to full energy level. u Increase from left to right: atoms become smaller, with greater nuclear charge. u Decrease as we go down a group.

Trends in Ionic Size u Cations form by losing electrons. u Cations are smaller that the atom they come from. u Metals form cations. u Cations of representative elements have noble gas configuration.

Ionic size u Anions form by gaining electrons. u Anions are bigger that the atom they come from. u Nonmetals form anions. u Anions of representative elements have noble gas configuration.

Configuration of Ions u Ions always have noble gas configuration. u Na is: 1s 2 2s 2 2p 6 3s 1 u Forms a 1+ 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 noble gas configuration. u They end up with the configuration of the noble gas after them.

Group trends u Adding energy level u Ions get bigger as you go down. Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

Periodic Trends u Across the period, nuclear charge increases so they get smaller. u 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 Iso electronic ions have the same # of electrons u Al 3+ Mg 2+ Na 1+ Ne F 1- O 2- and N 3- u all have 10 electrons u all have the configuration: 1s 2 2s 2 2p 6

Size of Isoelectronic ions u Positive ions that have more protons would be smaller. Al 3+ Mg 2+ Na 1+ Ne F 1- O 2- N 3-

Electronegativity u The tendency for an atom to attract electrons to itself when it is chemically combined with another element. u High electronegativity means it pulls the electron toward it. u Atoms with large negative electron affinity have larger electronegativity.

Group Trend u The further down a group, the farther the electron is away, and the more electrons an atom has. u More willing to share. u Low electronegativity.

Periodic Trend u Metals are at the left of the table. u They let their electrons go easily u Low electronegativity u At the right end are the nonmetals. u They want more electrons. u Try to take them away from others u High electronegativity.

Ionization energy, Electronegativity, and Electron Affinity INCREASE

Atomic size increases, shielding constant Ionic size increases