# Unit 6 – The Periodic Table

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Unit 6 – The Periodic Table

Origins of the Periodic Table
By the year 1700, only 13 elements had been identified Scientific discovery led to a higher rate of element discovery A logical organization of elements was needed for all the new elements

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 octaves Octave – repeating group of 8 elements

Mendeleev Dmitri Mendeleev (1869) arranged elements according to their properties Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, there was a repeating pattern to their properties This is known as Periodicity Mendeleev left some spaces on his table blank, but was able to predict the properties of the unknown elements

Mendeleev’s Periodic Table

Moseley Mendeleev’s table was imperfect – Te and I had to be reversed
Henry Moseley (1913) arranged elements according to atomic number The periodic repetition of chemical and physical properties when elements are arranged by atomic number is known as the Periodic Law

Modern Periodic Table The modern periodic table consists of Rows and Columns Rows - Horizontal Also known as Periods Numbered 1-7 Columns - Vertical Also known as Groups and Families Numbered 1-18

Classifying Elements The elements on the periodic table can be simply classified by groups Groups 1,2,13-18 (1A-8A) are known as the Representative Elements

Classifying Elements Groups of representative elements have the same valence electrons and Oxidation State Oxidation State is how many electrons are gained or lost by an atom in a chemical reaction Lost Electrons = Positive Oxidation State Gained Electrons = Negative Oxidation State Think of Oxidation State as the charge of the ion

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

Classifying Elements Groups 3-12 (3B-2B) , as well as the lanthanide and actinide series are known as Transition Metals

Metals The most common class of elements is Metals
Metals become cations What is a cation? How are they formed? Positively charged atom/positive oxidation state - Lose electrons Metals are generally solid (except Hg), conductive of heat and electricity, malleable, ductile, and shiny

Alkali Metals Group 1 elements are known as Alkali Metals
Alkali metals include Li, Na, K, Rb, Cs, Fr Alkali metals are generally dull, soft, and reactive – rarely found as free elements

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]7s1 How many valence electrons do all Alkali Metals have? One What is the oxidation state of Alkali Metals? +1

Alkaline Earth Metals Group 2 elements are known as Alkaline Earth Metals Alkaline earth metals include Be, Mg, Ca, Sr, Ba, and Ra Alkaline earth metals are harder, denser, and stronger than alkali metals Less reactive than alkali metals, but still rarely found as free elements

Alkaline Earth Metals Write the noble gas configuration for each Alkaline Earth Metal [He]2s2 [Ne]3s2 [Ar]4s2 [Kr]5s2 [Xe]6s2 [Rn]7s2 How many valence electrons do all Alkaline Earth Metals have? Two What is the oxidation state of Alkaline Earth Metals? +2

Transition Metals Elements in groups 3-12 (3B-2B) are known as Transition Metals Transition metals include Mn, Fe, Ag, Au, Mo, etc. Transition metals fill in the d orbital and often have multiple oxidation states Lanthanide and Actinide Series elements fill in the f orbitals – known as inner transition elements

Metalloids Elements that border the staircase on the periodic table are known as Metalloids Metalloids include: B, Si, Ge, As, Sb, Te, Po, At Metalloids have properties of both metals and nonmetals

Nonmetals Nonmetals are found to the right of the staircase on the periodic table Nonmetals generally become anions What is an Anion? How are they formed? Negatively charged atom/oxidation state - Gain electrons Nonmetals are often gases or dull, brittle solids Nonmetals generally show poor conductivity, ductility, and malleability

Halogens Group 17 elements are known as Halogens
Halogens include F, Cl, Br, and I Halogens are the most reactive nonmetals – often found in compounds

Halogens Write the noble gas configuration for each Halogen [He]2s22p5
[Ne] 3s23p5 [Ar] 4s23d104p5 [Kr] 5s24d105p5 How many valence electrons do all Halogens have? Oxidation State? Seven / -1 Why are the Halogens the most reactive non-metals? They are 1 electron short of having an octet.

Noble Gases Elements in group 18 are known as Noble Gases
Noble Gases include He, Ne, Ar, Kr, Xe, Rn Noble gases are extremely unreactive

Noble Gases How many valence electrons do all Noble Gases have? Eight
Write the electron configuration for each Noble Gas 1s2 [He]2s22p6 [Ne]3s23p6 [Ar]4s23d104p6 [Kr]5s24d105p6 [Xe]6s25d106p6 How many valence electrons do all Noble Gases have? Eight Why are Noble Gases so unreactive? They contain a full octet – atoms gain/lose electrons to achieve noble gas notation

Other Groups All other groups can be identified by the top most element in that group. Ex: Group 15 can be called the Nitrogen Group Oxidation State: -3 Q: What is another name for Group 16? A: Oxygen group Q: Oxidation State A: -2

Periodic Trends The elements on the periodic table show repeating trends related to electron configuration

What is the trend for Oxidation State?

Atomic Radius The Atomic Radius is ½ the distance between nuclei of bonded atoms from the same element Atomic radius decreases from left to right across a period Atomic radius increases from top to bottom in a period

Why? Not changing energy level, but increasing nuclear force (more positive charge in nucleus)

Ionization Energy If an atom is becoming an ion, it is gaining or losing electrons in an effort to have an octet (8 valence electrons)

Ionization Energy The energy required to remove an electron from an atom is called Ionization Energy

Ionization Energy 1st Ionization Energy- energy required to remove 1st electron from an atom 2nd Ionization Energy- energy required to remove 2nd electron from an atom 2nd Ionization Energy is ALWAYS higher than the 1st 3rd Ionization Energy- energy required to remove 3rd electron from an atom 3rd Ionization Energy is ALWAYS higher than the 1st or 2nd

Ionization Energy IE Decreases as you move down a group Why?
Electron is further away

Ionization Energy IE Increases as you move across a period Why?
You are in the same energy level but have more nuclear charge

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

Ionization Energy Write the electron configuration for Be 1s22s2
How many valence electrons does Be have? 2 Why 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.

Ionization Energy Move across the period. Write the electron configuration for F. 1s22s22p5 How many valence electrons does F have? 7 Why 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

Ionization Energy

Electron Affinity Electron 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.

Ionic Size Cations are smaller than the atoms from which they form (less electrons) Anions are larger than the atoms from which they form (more electrons)

Ionic Size Across the period, nuclear charge increases so they get smaller. Energy level changes between anions and cations. N3- O2- F1- B3+ Li1+ C4+ Be2+

Electronegativity Electronegativity is the ability for an atom to attract electrons in a compound Electronegativity increases from left to right in a period Electronegativity decreases from top to bottom in a group

Electronegativity We do not consider noble gases when talking about electronegativity because they do not bond. What is the most electronegative element? Fluorine

Electronegativity Write the electron configuration for Li 1s22s1
How many valence electrons does Li have? 1 Why 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

Electronegativity Move across the period. Write the electron configuration for O. 1s22s22p4 How many valence electrons does O have? 6 Why 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.

Electronegativity

Ionization energy, Electronegativity, and Electron Affinity INCREASE

Atomic size increases, Ionic size increases