History and Trends of the Periodic Table

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Year of Discovery

Ch.8:2,4,5,8,19,34,35,37,51,56,61,68

Johann Dobereiner – Noticed several “triads” of similarly behaving elements (1829)

His practice of grouping elements led to the concept of periodicity (The repeating of chemical properties on a sequential basis)

Element Grouping

Element OrganizationJohn Newlands – Noticed periodicity based on elemental mass

and sorted elements by mass (1864)•Law of Octaves – every 8th (known) element was similar•His contribution wasn’t recognized till years later

Dmitri Mendeleev (1869) “Father of the periodic table”

•Also arranged elements by mass and property, but included empty space for irregular jumps in mass

•These accounted for many undiscovered elements successfully predicted later

Modern Periodic Law

Henry Moseley (Rutherford’s assistant)• Discovered a method for counting each element’s protons with x-rays•Found that if elements were organized by atomic number, and not mass, the problems disappeared. (1912)

Periodic law – Properties vary with their atomic numbers in a periodic way. (Current basis of P.T. organization)

•Problems still arose as new elements were discovered and didn’t always fit nicely when ordered by mass

Periodic Elements

Atomic Number - Number of protons• Basis of Periodic Table organization.• Distinguishing attribute of each element (always an

integer)

•Name - full name of element

•Symbol - 1,2, or 3 letter representative symbol

•Atomic Mass - weighted average mass of relative isotopes•Atomic mass units (amu) or grams/mole

Groups (families) - Vertical columns • Share chemical traits (properties)• IUPAC - numbered groups 1 - 18• NACPT notation - 1A - 8A / 1B - 8B

Periods (series) - Horizontal rows• Indicates highest energy electron level n (1-7)

Other Divisions: – Representative elements - (1A - 8A) (s and p block)• Consistently "periodic" based on valence electrons

– Transition metals - (1B - 8B) (d and f block)• Many exceptions exist for electron configurations• Characteristics are less periodic• f-block (inner-transition metals) (Lanthanide and Actinide series)

Classification of the ElementsIUPAC: 1-18

NACPT: 1A-8A1B-3B

*Noble gases are part of the Representative elements

Valence electrons are the outer shell electrons. The valence electrons are the electrons that take part in chemical bonding.

1A 1ns1

2A 2ns2

3A 3ns2np1

4A 4ns2np2

5A 5ns2np3

6A 6ns2np4

7A 7ns2np5

Group # of valence e-e- configuration

Core electrons: all non-valence electrons

Octet (8) Rule• Many chemical properties are determined by an

atom’s valence electrons

• Atoms seek to obtain an electron configuration like that of a Noble Gas

• A full outer electron shell has 8 valence electrons “oct”-et (of representative elements)

• Helium reaches full valence with only 2 electrons

Neon1s2 2s2 2p6

+1

+2

+3

-1-2-3

Cations and Anions Of Representative Elements

Metals (~75% of elements)

•Left and middle of periodic table

•Solid at RT (except Hg)

•Lustrous (shiny)

•Malleable (flattened out)

•Ductile (drawn to strips)

•High melting points (~950 – 3,700 °C)

•Good conductors of heat and electricity

King of Random: The Metal Melterwww.youtube.com/watch?v=GCrqLlz8Ee0

Nonmetals (~18% of elements)

• Top Right corner of periodic table + Hydrogen• Typically gaseous or soft, crumbly solids at room temperature

(low melting temps).

• Exceptions: Bromine(l) and Cdiamond (s)

• Poor electrical conductors (Carbon is a thermal conductor)

• Combine with other non-metals to form molecules.

•Combine with metals to form Ionic compounds

Discovery of the 6 Noble Gases (1898 – 1900)

Sir William Ramsay

• Difficult to detect due to inert nature• Ramsay reacted N2(g) with Mg(s) to form Mg3N2(s) along with a volume of unknown gas that would not react.

• Using a discharge tube, they noted the emission spectrum was unique.• It was called Argon “the lazy one”

Metalloids (~7% of elements)• Adjacent the Stair-step line that separates metals

from non-metals

• Have shared properties of both metals and non metals

• Metallic luster

• Brittle or crumbly

• Fair conductors

Trans-Uranium Elements– Only elements up to atomic 92 (Uranium) have been found naturally occurring (trace amounts of Pu & Np)– Nuclear reactions have produced the remaining elements synthetically

Radioactive ElementsAll elements after #83, starting with Polonium, are radioactive

(element decays rapidly)Crash Course: The Periodic Table

www.youtube.com/watch?v=0RRVV4Diomg

Periodic trend summary

• Effective Nuclear charge – nuclear pull “felt” by electrons

•Atomic radii - Distance between nucleus and outer e-

• Ionic radii - same as atomic radii, but distance for each atom's common ion

• Ionization energy - energy for atom to lose an electron (lower means more likely to lose an electron)

• Electron affinity - energy released when an atom gains an electron (higher means more likely to gain electron)

• Electronegativity - measure of each atom's attraction towards bonding electrons in a molecule

Effective nuclear charge (Zeff) is the attractive force from the nucleus felt by an electron with all forces taken into account.

Shielding effect: Repulsive forces from other electrons that lessen the net force felt from the positive nucleus

Harder to remove, b/c no shielding effects (e-/e- repulsions)

Effective Nuclear Charge (Zeff Total – Core electrons)

Increasing Zeff

decr

easi

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eff

11Na

12Mg

13Al

11

12

13

10

10

10

1

2

3

186

160

143

ZeffCore (e-)Z Radius (pm)Additional P+ have greater pull than

additional valence e-

Greater Zeff leads to smaller atomic radius

Trends in Atomic Radii

Atomic radii get smaller as atoms get more massive across a period

Atomic radii increases going down group (adding n)Period:

2

3

5

4

6

Only representative elements shown

Zeff = 5.76 Zeff = 3.14Zeff = 1.28

Per n, adding Protons has greater effect than adding valence e-

Decreasing Atomic radii

Cation is always smaller than atom from which it is formed. Same # of protons (+) pulling in less electrons (-)

Anion is always larger than atom from which it is formed. More electrons (-) being held by Same # of protons (+)

(152 pm) (72 pm)

(133 pm)

(78 pm)

25

Comparison of Atomic Radii with Ionic Radii

Cations get smaller Anions get larger

The Radii (in pm) of Ions of Familiar Elements1 m = 1 x 1012

pm

Filled n=1 shell

Filled n=2 shell

Filled n=3 shell

Filled n=4 shellFilled n=5 shell

Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. (Cation forming)

Noble gases have full valence orbitals = very stable = unreactive

Alkali metals have 1 valence e- and easily give it up

General Trends in First Ionization Energies

Increasing First Ionization Energy

Dec

reas

ing

Fir

st I

oniz

atio

n E

nerg

y

Very unlikely to give away electrons

Very Likely to give away electrons

Higher energy levels electrons can more readily leave (further from nucleus, more core electrons shielding ~ lower Zeff)

I1 < I2 < I3 < …

“Typical” last Ionization:

Note the large jump in energy afterward

Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.

X (g) + e- X-(g)

F (g) + e- F-(g)

O (g) + e- O-(g)

EA = +328 kJ/mol

EA = +141 kJ/mol

Larger release of energy indicates more stable anion (more likely to gain electron)

Why would Alkali metals have a higher affinity than Alkaline Earth metals?Why is Nitrogen 0? Half-full p-subshell is more stable than 4/6

Halogens (17/7A) most likely to gain one electron

Electronegativity - attraction of shared electron cloud Linus Pauling mathematically determined the measure of attraction between an atom's nucleus and its valence electrons by analyzing bond strength of various compounds.

No values because they

don't form compounds

Periodic trend summary

• Effective Nuclear charge – nuclear pull “felt” by electrons

•Atomic radii - Distance between nucleus and outer e-

• Ionic radii - same as atomic radii, but distance for each atom's common ion

• Ionization energy - energy for atom to lose an electron (lower means more likely to lose an electron)

• Electron affinity - energy released when an atom gains an electron (higher means more likely to gain electron)

• Electronegativity - measure of each atom's attraction towards bonding electrons in a molecule

1. a) Place the neutral atoms in order from smallest to largest Radius: N, Na, P

b) Explain what is causing the change across a period (L to R).

2. a) Place the neutral atoms in order from smallest to largest Ionization energy: K, Se, Rb

b) If an atom has a low value for Ionization energy, what action is this specifically referring to and how likely will it happen?

3. The first six Ionization energies (kJ/mol) are listed sequentially for Aluminum. Make two observations from the data.

Practice Questions

4. a) Place the atoms in order from smallest to largest electron affinity energy: Ca, Se, Kr

b) If an atom has a low value for electron affinity, what action is this specifically referring to and how likely will it happen?

5. What happens to the relative atomic radius of neutral Potassium atoms and Bromine atoms after they ionize and combine to form KBr?

6. Place the elements in order of increasing electronegativity: F, As, N, Ne

Practice Questions

Descriptive Chemistry

•Study of the elements and the compounds they form.

•Physical and Chemical Properties

•Similar for each group/family

Hydrogen

• Lightest and most abundant element• Non metal (though displayed in 1A) • Considered a family of its own• Colorless, odorless, and tasteless gas• Has chemical properties similar to both alkali

metals (reactivity) and halogens (physically).

• Occurs "di-atomically" H2, not H1s1

The Alkali Metals Group 1A Elements (ns1)

M M+1 + 1e- (loss of 1 valence e-) 2M(s) + 2H2O(l) 2MOH(aq) + H2(g) (Hydrogen displacement)

4M(s) + O2(g) 2M2O(s) (Combustion/Oxide formation)

• Very chemically reactive• Form +1 charge cations• Conductive and Lustrous• Soft at room temperature• React with water

• Lithium batteries • 133Cs in atomic clocks

Incr

easi

ng r

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ivit

y

The Alkali Metals: Group 1A Elements (ns1)

• Na+ and K+ : mediate conduction across membrane synapses for

nervous system

Cs in slow-mo

Alkaline-Earth Metals/Group 2A Elements (ns2)

Luminescent & radioactive

M M+2 + 2e-

M(s) + 2H2O(l) M(OH)2(aq) + H2(g) (M = Mg, Ca, Sr, or Ba)

Alkaline-Earth Metals• Solid at room temperature• Forms +2 charge cations• Reactive at higher periods

• Milk of Magnesia (Mg(OH)2)

• Strontium: flares, red fireworks• Barium: rat poison,

gastrointestinal x-ray, green in fireworks

Boron Family: Group 3A Elements (ns2np1)Tend to form +3 ions

Boron: found in lab glasswareAluminum: very common in alloys (light weight)

Periodic Videos:Gallium

Carbon Family: Group 4A Elements (ns2np2)• Typically do not form ionic bonds, but covalent • Carbon: organic chemistry: vitamins/drugs. Main

component of all biomolecules (protein, fat, sugar, DNA)• Silicon: heavily used in electronics

Nitrogen Family: Group 5A Elements (ns2np3)• Tend to form -3 ions• Nitrogen: most abundant atmospheric gas, found in all

proteins and DNA, very inert as N2

• Phosphorous: Very reactive, match heads; present in DNA, ATP, and lipids)

Periodic Videos:Phosphorous

Bi oxide

Oxygen Family: Group 6A Elements (ns2np4)•Tend to form -2 anions

• Almost all life is sustained by aerobic respiration which requires oxygen (electron carrier)• O2 is required for combustion (fire)

• O forms many important compounds (oxides) (CO2, NO3, PO4, SO4)• Sulfur: pure form has distinct smell of rotten eggs, found in 2 amino acids

Burning sulfur

Barking Dog

Properties of Oxides Across a Period

basic acidic

Oxygen can form compounds with elements of

various groups

59

Group 17/7A Elements (ns2np5): Halogens

Group 17/7A Elements (ns2np5): Halogens

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• Pure forms are diatomic (ex. F2, Cl2)

• Form salts when they react with metals• Large electronegativities; Highest e- affinity• Very reactive• Bleach contains chlorine compounds• Fluoride: toothpaste• Iodine used as disinfectant

Periodic Videos:Chlorine

Noble Gases: Group 18/8A Elements (ns2 np6)Completely filled ns and np subshells.Inert: Highest ionization energy/low e- affinities

Low tendency to lose/accept electrons.Colorless, odorless, and tasteless gasesLow boiling/ freezing points

Compounds of the Noble Gases

A number of xenon compounds exist: XeF4, XeO3, XeO4.

A few krypton compounds have been prepared, such as KrF2.

PtF6

gas

Transition Metals

Zinc present in many DNA unwinding proteins

Fe binds O2 in Hemoglobin

Necessary in trace amounts in living organisms

Periodic Videos:Zinc

Found to be “less periodic”– Difficult to purify and thus harder identify– Did not fit into major groups, overall similar properties– Contain the Coinage and Precious metals (Au, Ag, Pt)

Inner Transition MetalsAlso called the Lanthanide and Actinide Series• Paramagnetic• Actinides are radioactive and most are synthetically made

Uranium:U-235 used in

nuclear reactors

U

Neodymium

Nd

Nd2Fe14B alloy