12/2/14 Today I will discuss the development of the periodic table Warm Up – Name two ways we have used the periodic table.

CP Chemistry Chapters 5/6 The Periodic Table CP Chemistry Chapters 5/6

Periodic Table To organize the elements! Why did scientists develop the periodic table? To organize the elements! The periodic table worked so well that scientists were actually able to predict elements that had not been discovered yet!

Forerunners 1700’s – only about 30 elements known 1800’s – about 60 elements known Scientists needed a good way to classify the elements

Periodic Table Using either the book, or your “smart” devices, you have 20 minutes to answer the questions regarding the scientists of the periodic table. Ready, set, go!

J.W. Dobereiner – early 1800’s Triads – groups of three elements with similar chemical properties Li, Na, K Ca, Sr, Ba Cl, Br, I The middle element has approximately the average mass and average density of the other two!

J.A.R. Newlands - 1865 Law of Octaves – when elements are arranged by increasing atomic mass, property patterns repeated every eight elements. 1st element like the ninth 2nd element like the tenth, etc. No one believed him because of the relationship between chemistry and music!

Dmitri Mendeleev - 1869 The Father of the Periodic Table! Placed elements on index cards and arranged them based on repeating patterns of properties Increasing Atomic Mass Greatest accomplishment was leaving blank spaces for elements not yet discovered!

Dmitri Mendeleev The Father of the Periodic Table!

Henry Moseley – early 1900’s Recognized that the patterns are based on atomic number, not atomic mass. Fixed a few issues with Mendeleev’s table. Ar (At# 18, AtW 39.95) K (At# 19, AtW 39.01) Periodic Law – properties of elements are periodic functions of their atomic number. Cycles are observed as the Atomic Number increases

Homework 5-1 R&R (Front only)

12/3/14 Objective – identify the organization of the periodic table Warm up – What did Newlands contribute to the understanding of the elements?

Periodic Table Activity Using your smart devices or the book, complete the periodic table activity

12/4/14 Objective – identify the organization of the periodic table Warm up – Give 2 examples of alkaline earth metals.

Reading the Periodic Table Modern Periodic Table - Arrangement of elements in order of their atomic number so that elements with similar properties fall into the same column. Groups or Families – vertical columns Periods – Horizontal rows of elements

“s” n= 1 2 3 4 5 6 7 “p” “d” “f”

“s” & “p” = main-group elements “d” = transition metals “f” = inner transition metals

Periodic Blocks What do the blocks of the periodic table tell us?

Group Numbers 1 18 1A 8A 2 13 14 15 16 17 2A 3A 4A 5A 6A 7A 3 4 5 6 7 9 10 11 12 3B 4B 5B 6B 7B 8B 1B 2B

Family Names Group 1 – Alkali metals Group 17 – Halogens Group 2 – Alkaline-earth metals Group 17 – Halogens Group 18 – Noble gases Groups 13-16 1 2 3 4 5 6 7 8 9 10 11 12 18 13 14 15 16 17

Lanthanides Actinides Groups 3-12 - Transition Metals Below the Table - Inner Transition Metals Lanthanides Actinides 1 2 3 4 5 6 7 8 9 10 11 12 18 13 14 15 16 17

Family Properties Alkali metals – very malleable, soft metals, very reactive Alkaline Earth metals – malleable, harder metals, reactive Halogens – gaseous (F, Cl, Br) or solid (I, At) nonmetals, often toxic, very reactive Nobel gases – inert gases

=metaloid (semi-metal) Metals/Nonmetals Non-Metals Metals

Properties of Metals/Nonmetals Luster or shine Good conductors of heat and electricity Typically solid at room temp (except Hg) Malleable – can be hammered into sheets Ductile – can be drawn into wire

Properties of Metals/Nonmetals Typically dull finish Poor conductors of heat and electricity Variety of forms at room temp (solid, liquid, gas) Neither malleable nor ductile Metalloids (semi-metals) Properties of both metals and non-metals Useful in computers

Families The electrons in their outermost shells are the same! Why do elements in the same family have similar properties? Let’s look at the electron configurations H – 1s1 Li – 1s22s1 Na – 1s22s22p63s1 K – 1s22s22p63s23p64s1 Rb - 1s22s22p63s23p64s23d104p65s1 The electrons in their outermost shells are the same!

Families Why is that important? Valence electrons – electrons in the outermost shell responsible for an atom’s chemical behavior and properties.

Homework 5-2 R&R

12/5/14 Today I will determine the periodic trends for atomic radius and ionic radius Warm Up – How many valence electrons do the halogens have? Hint: Look at the electron configuration for F & Cl.

Periodic Trends Periodic Trends- properties that change in predictable ways as you move through the periodic table.

Atomic Radius Atomic Radius – The distance from the center of an atom’s nucleus to its outermost electron (half the distance between two neighboring nuclei)

Atomic Radius Atoms get larger going down a column New energy levels are added Atoms get smaller moving across a period ???? – Aren’t we adding electrons? We are also adding protons which pull the electrons in tighter!

Ionic Radius Ionic Radius – Atomic radius of an atom that has gained or lost an electron Where do ions come from?

Ionic Radius The electrons in their outermost shells are the same! Remember Valence Electrons: Let’s look at the electron configurations H – 1s1 Li – 1s22s1 Na – 1s22s22p63s1 K – 1s22s22p63s23p64s1 Rb - 1s22s22p63s23p64s23d104p65s1 The electrons in their outermost shells are the same!

Valence Electrons Elements of the same family also have the same number of valence electrons (outermost shell) 1 8 2 3 4 5 6 7

Ionic Charge – common charge Octet Rule – Atoms will lose or gain electrons to have a full set of valence electrons All atoms want to have 8 electrons in their outermost shells (except only 2 in the first shell). When they do they have the same configuration as the closest noble gas! Noble gas configurations are very stable!

Ionic Charge – common charge Elements of the same family tend to form the same ions! +1 +2 +3 ±4 -3 -2 -1

Ionic Radius How does gaining or losing an electron affect the atomic radius? Cations – positive ions that have lost electrons Smaller than the neutral counterpart because they have less electrons to pull in, so they can be pulled tighter! Anions – negative ions that have gained electrons Larger than the neutral counter part because they have more electrons to spread out!

Ionic Radius Cations (smaller) tend to occur to the left of the periodic table Anions (larger) tend to occur to the right of the periodic table

Homework Atomic/Ionic Radius WS

12/8/14 Today I will name and explain trends that occur on the periodic table. Warm Up – Which atom in each pair is larger? a) Al or Si b) Ge or Sn c) Ba or Pb d) Mg or its most common ion e) S or its most common ion

Valence Electrons Elements of the same family also have the same number of valence electrons (outermost shell) 1 8 2 3 4 5 6 7 2

First Ionization Energy Ionization Energy – the energy needed to remove an electron The more tightly bound the electrons, the harder they are to remove. Greater as you move across More protons provide more pull Smaller as you move down Added energy levels because they are further from the protons’ pull Electron shielding – Inner electrons shield outer ones from the pull of the nucleus

Successive Ionization Energy Successive Ionization Energy – the energy needed to remove another electron after the first Greater because of the greater number of protons than electrons Much greater jump as you move to closer energy levels

Valence Electrons Mg – Why would the third ionization be much larger than the second?

Electronegativity Electronegativity – the ability of an atom to attract an electron All atoms want to have 8 electrons in their outermost shells (except only 2 in the first shell).

Electronegativity Which elements want to gain electrons? 1 8 2 3 4 5 6 7 2

Electronegativity Higher as you move across elements are more likely to gain an electron due to valence number. (until noble gases!) Electronegativity decreases down a group easier for the nucleus to pull in an electron that is closer (in a lower energy level).

Electron Affinity Electron Affinity - the energy change that occurs when an electron is acquired by a neutral atom. Related to electronegativity because the more electronegative an atom, the more energy it will aquire Follows the same trend as electronegativity

Trend Review Atomic Radius Decreases Ionization Energy Increases Electronegativity/Electron Affinity Increases Atomic Radius Increases Ionization Energy Decreases Electronegativity/Electron Affinity Decreases

Trend Review Cations on the left are generally smaller Anions on the right are generally larger

Special Cases/Exceptions Hydrogen – hydrogen resides in group 1, but it’s not a metal! It is a very different element. First energy level – We said that all elements want 8 electrons in the outermost shell, but there is an exception. The first energy level, only wants two electrons! (Hydrogen and Helium) This is why helium (with only 2 valence) is a noble gas!

Homework Ionization energy, Electronegativity & Electron Affinity WS Graphing Project starting tomorrow

12/9/14 Today I will graph and analyze periodic trends Warm Up – Explain why the first ionization energy trend is what it is…

12/10/14 Today I will graph and analyze periodic trends Warm Up – Explain the jump in ionization energies between a nobel gas and the following alkali metal.

12/11/14 Today I will graph and analyze periodic trends Warm Up – Sketch the periodic table and give an overview of each of the trends

12/12/14 Today I will write noble gas configurations Warm Up – Write the electron configuration for Ge.

Noble Gas Configuration Let’s look at the electron configurations again H – 1s1 Li – 1s22s1 Na – 1s22s22p63s1 K – 1s22s22p63s23p64s1 Rb - 1s22s22p63s23p64s23d104p65s1 Everything up to the last numbers is always the same! Is there a shortcut we can take?

Noble Gas Configuration 1. Place the name of the noble gas before the element in brackets, and 2. continue writing the electron configuration. Na – 1s22s22p63s1 [Ne]3s1 Ge - 1s22s22p63s23p64s23d104p2 [Ar] 4s23d104p2

Writing in Short-hand notation Write the short hand notation for the following: V Rb I Hg U W [Ar] 4s2 3d3 [Kr] 5s1 [Kr] 5s2 4d10 5p5 [Xe] 6s2 4f14 5d10 [Rn] 7s2 5f4 [Xe] 6s2 4f14 5d4

Homework Finish the graphing project – due Monday!

12/15/14 Today I will review periodic trends Warm Up – Why do metals tend to form positive ions while non-metals tend to form negative ions?

Periodic Trends Review 1. Periodic Law states that elements are a periodic function of their atomic #. This means that you see repeating patterns when elements are lined up by atomic #. This is important because it’s the basis of the modern periodic table!

Periodic Trends WS

12/16/14 Today I will review chapter 5/6 Warm Up – Write the noble gas configurations for Se, Nb, and Sm.

12/17/14 Today I will review for the chapter 5/6 exam Warm Up – List and describe the 4 people responsible for the development of the modern periodic table.