Unit 5 – 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 triadsTriad – three elements with similar

properties (ex: Cl, Br, I)

J.R. Newlands (1864) organized elements in octavesOctave – 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 PeriodsNumbered 1-7

Columns -VerticalAlso known as Groups and FamiliesNumbered 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 cationsWhat 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 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 / -1Why 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 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 indentified by the top most element in that group.Ex: Group 15 can be called the Nitrogen

GroupOxidation State: -3Q: What is another name for Group 16?A: Oxygen groupQ: Oxidation StateA: -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 atom2nd Ionization Energy is ALWAYS higher

than the 1st

3rd Ionization Energy- energy required to remove 3rd electron from an atom3rd 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.

Li1+

Be2+

B3+

C4+

N3-O2- F1-

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