Unit 5 The Periodic Table The how and why

Newlands u Arranged known elements according to properties & order of increasing atomic mass u Law of Octaves – pattern of chemical & physical properties repeated every 8 elements

Mendeleev u Created 1 st periodic table (63 elements) u Ordered by increasing atomic mass u Predicted pattern of missing elements u Started new rows and lined up columns to organize elements with similar properties u Rearranged elements so similar properties would line up correctly

The Modern Table u Moseley- determined the atomic number for each known element. u Elements are still grouped by properties u Similar properties are in the same column u Ordered by increasing atomic number u Added a column of elements Mendeleev didn’t know about – noble gases

Periodic Law u When elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals

u Horizontal rows are called periods u There are 7 periods

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

1A2A 3A4A5A6A7A 8A 3B4B5B6B7B8B 1B2B IAIIA IIIBIVB VB VIBVIIB VIIIB IIIAIVAVA VIAVIIA VIIIAIBIIB Other Systems

1A 2A3A4A5A6A 7A 8A 0 u The elements in the A groups are called the representative elements

Transition metals l The Group B elements

u These are called the inner transition elements and they belong here

Three Classes of Elements u Metals u Nonmetals u Metalloids

Metals

l Ductile – drawn into wires l Malleable – hammered into sheets l All solid at room temperature (except Hg- Mercury) l Conductors of heat and electricity l Families –1 - Alkali –2 - Alkaline Earth –Transition (B groups)

u Group 1A are the alkali metals u VERY reactive because one valence e - Found as compounds in nature Not including H!

u Group 2A are the alkaline earth metals u Still highly reactive but not as much so u as alkali metals (2 valence e-)

Transition Metals u The weird ones… u May lose different #s of valence electrons depending on the element with which it reacts u Less reactive than alkali or alkaline earth metals u Good conductors of electricity & heat, ductile, malleable

Inner Transition Metals u 1 st row = lanthanides Shiny metals similar in reactivity to alkaline earth metals u 2 nd row = actinides Unstable nuclei – all radioactive

Non-metals

l Most are gases, some solid, and 1 liquid (Br) l More variation than metals l Families –Halogens (Group 17 or 7A) –Noble Gases (Group 18 or 8A)

u Group 7A is called the Halogens u Most reactive non-metals – 7 valence React frequently with alkali metals

Group 8A are the noble gases Low reactivity, very stable, inert

Metalloids or Semimetals

Metalloids l Border the staircase between metals and nonmetals l Properties – similar to metals and nonmetals

Part 2 Periodic trends Identifying the patterns

What we will investigate u Atomic size how big the atoms are u Ionization energy How much energy to remove an electron u Electronegativity The attraction for the electron in a compound

What we will look for u Periodic trends How those things vary as you go across a period u Group trends How those things vary as you go down a group u Why? Explain why these variations exist

Atomic Size u Where do you start measuring? u The electron cloud doesn’t have a definite edge. u Scientists focused first on diatomic elements -- measured more than 1 atom at a time

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

Atomic Size - Periodic Trends u The positive nucleus pulls on electrons u Periodic trend As you move across a period, elements have more protons The charge on the nucleus gets bigger The outermost electrons of each element are in the same energy level So there is more pull on the outermost electrons as you move across

Periodic Trends u As you go across a period, the radius gets smaller. u Same outermost energy level u More nuclear charge u Pulls outermost electrons closer NaMgAlSiPSClAr

Atomic Size – Group Trends u The positive nucleus pulls on electrons u Group Trend As you go down a group, you add energy levels Outermost electrons not as attracted by the nucleus

+ Shielding u Increasing numbers of electrons between the nucleus and the valence electrons tends to decrease the force between the nucleus & the valence electrons

+ Shielding u The electron on the outside energy level has to look through all the other energy levels to see the nucleus

Shielding u The electron on the outside energy level has to look through all the other energy levels to see the nucleus u A second electron has the same shielding u In the same energy level (period) shielding is the same +

Shielding u As the energy levels changes the shielding changes u Moving down the group More energy levels More shielding Outer electron less attracted + No shieldingOne shieldTwo shieldsThree shields

Group trends u As we go down a group Each atom has another energy level More shielding The atoms get bigger H Li Na K Rb

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

Atomic size increases,

IONIZATION ENERGY

It’s all about stability u Alkali metals are more stable if they lose an electron u Example Sodium ([Ne] 3s 1 ) Getting rid of the 3s 1 electron makes sodium more stable and creates a sodium ion (Na 1+ )

Ionization Energy u The amount of energy required to completely remove an electron from a neutral atom. u The energy required for the 1 st electron is called the first ionization energy

Ionization Energy u The 2 nd ionization energy is the energy required to remove the second electron u Always greater than 1 st IE u The 3 rd IE is the energy required to remove a third electron u Greater than 1 st or 2 nd IE

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

Group trends u As you go down a group first IE decreases Valence e - farther from nucleus More shielding

Periodic trends u First IE increases from left to right across a period Increased nuclear charge from added proton Electron shielding not an issue b/c valence are all in same energy level u Exceptions at full and 1/2 full orbitals Lower IE b/c offer stability to atom

Ionization energy INCREASE

How to remember? HI LO

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 l Li has lower IE than H l more shielding l outweighs greater nuclear charge Li

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

First Ionization energy Atomic number H He l B has lower IE than Be l same shielding l greater nuclear charge l By removing an electron we make s orbital full 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 gets to 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

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

First Ionization energy Atomic number

Electronegativity

u There’s an electron tug of war between atoms in a compound u The tendency for an atom to attract electrons to itself when it is chemically combined with another element u How “greedy” u Large electronegativity means the atom pulls the electron towards itself

Group Trend u As you move down a group More shielding Less attraction for electrons Lower electronegativity

Periodic Trend u As you move across a period from left to right, Nuclear charge increases Greater electronegativity

Electronegativity INCREASE

How to remember? HI LO

All 3 trends HI LO