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The Periodic Law
5-1 History of the Periodic Table
Objective: Covers the work of Mendeleev and other chemists in developing the
periodic table and explains how the periodic law is used to predict elements’
physical and chemical properties.
Mendeleev and Chemical Periodicity
 Dmitri Mendeleev noticed that when elements were arranged in order of
increasing atomic mass, certain similarities in their chemical properties appear in
regular intervals.
o Periodic refers to a repeating pattern.
 Mendeleev grouped elements with similar properties.
 Mendeleev left gaps in his periodic table for elements that were not known based
on atomic masses and properties.
o He named one these unknown elements, ekasilicon, which was later
isolated and named germanium.
 *Mendeleev arranged his table by increasing atomic mass.
Moseley and the Periodic Law
 Discovered the concept of atomic number (based this information on charges).
 He believed that the periodic table was organized by atomic number and not
atomic mass, which is was correct.
 Periodic Law – the physical and chemical properties of the elements are periodic
functions of their atomic number.
The Modern Periodic Table
 Noble Gases, Lanthanides, and Actinides: Know where it is located, what its
name means, but I don’t care if you memorize the history and scientists.
5-2 Electron Configuration and the Periodic Table
Objective: Explains the relationship between electron configuration and the
arrangement of elements in groups, blocks, and periods of the periodic table, as
well as the elements’ general properties.


The stability of an atom depends upon the electron configuration.
Usually the outer electrons determine chemical reactivity.
Periods and Blocks of the Periodic Table
 Groups (Families) – vertical columns
o Elements in groups have similar properties.
 Periods – horizontal rows of the periodic table.
 Fig. 5-5 pg. 129
The s-Block Elements: Groups 1 and 2 (IA and IIA)
 All are chemically reactive metals.
 Outermost electrons of these atoms are in the s orbital.
 Alkali Metals – The elements of group 1 (IA)
o Have only 1 valence electron
o All metals have silver appearance and are soft.
o They are not found free in nature because they are so reactive.
o React vigorously with non-metals (esp. halogens)
o Are extremely reactive with water.
 Most be stored in kerosene because they can react with the little
water vapor that is in the air.
o All have low melting points.
 Alkaline-Earth Metals – elements of groups 2 (IIA)
o Have 2 valence electrons.
o Are very reactive, and not found in nature.
 Are less reactive than alkali metals.
o Are harder, denser and stronger than alkali metals.
p-Block Elements: Group 13-18 (IIIA to VIIIA)
 Outermost electrons of these atoms are in the p orbital.
 Boron Group (IIIA)
o Have 3 valence electrons
o Form 3+ ions.
o Aluminum: Al is the 3rd most abundant element in the earth’s crust.
 Common, inexpensive metal.
 Does not corrode.
 Used in structural materials, cans, pots and pons.
 Carbon Group (IVA)
o Carbon: found in limestone (CaCO3)
 Carbon and compounds make up fossil fuels.
 Exists as diamond and graphite.
 Difference between types depends upon bonding.
 Compounds of H and C are called hydrocarbons.
 Make petroleum, coal and natural gas.
o When burned form CO and CO2
 Common pollutants.
 Make plastics and medicine.
 The study or carbon compounds is called organic chemistry.
o Silicon: 2nd most abundant element in earth’s crust.
 SiO2 is called silica (sand).
 Major component of glass.
o Transition metals are sometimes added to glass to
give color.
 Ex. Cobalt gives blue color.
 Silicates are compounds that contain Si and O and other elements.
 The structural material of most rocks.
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 Is used as a semiconductor, computer chips, and solar cells.
o Germanium: used to make electronic devices.
o Tin: used to coat items, such as “tin cans”.
o Lead: used in automobiles, lead storage batteries, and as a structural
material.
Nitrogen Group (VA)
o Nitrogen: Makes up about 80% of the earth’s atmosphere.
 Nitrogen is essential for life, but most organisms cannot get it from
the atmosphere. Organisms rely on nitrogen compounds.
 Bacteria converts nitrogen gas into nitrogen compounds.
o Nitrogen fixation.
 It is used in fertilizers, ammonia and detergents.
o Phosphorous: Combines with oxygen to form phosphates (common in
minerals)
 Major use is in phosphoric acid.
 Used in fertilizers and detergents, but also in soft drinks to
give a tart taste.
Oxygen Group (VIA)
o Oxygen: most abundant element on Earth
 Makes up 23% of atmosphere and 46% (by mass) of earth’s crust.
 Combines with elements to form oxides.
 Is required for respiration and is involved in combustion.
 Elemental oxygen exists in a diatomic form (O2) and in Ozone (O3)
o Sulfur: most common characteristic is the unpleasant odor.
 Largest use of sulfur is in sulfuric acid, used in batteries.
 Used in paints, lubricants, plastics, explosives, …
 Sulfur is common in compounds including fools gold, pyrite.
Halogens (VIIA)
o All of the halogens found in elemental form are diatomic (2-atoms)
 Ex. Ex. Cl2, F2, I2, …
 7 at 7
o React with most metals and many nonmetals.
o They are highly reactive and do not occur free in nature.
o They all have a strong attraction for electrons.
 Fluorine: Is the most reactive of all the elements.
 Found in CFC’s (chlorofluorocarbons)
o Used in refrigerators, air-conditioners, and aerosols.
 Harmful to ozone.
 Chlorine: occurs naturally in salt.
 Is added to water as a disinfectant.
 Used in bleach and PVC plastics.
 Bromine: used in pesticides, photographic film and fire retardants.
 Iodine: used as an antiseptic.
Noble Gas (VIIIA)
o They are all non-reactive elements.
o All have 8 valence electrons and don’t want to loose or gain electrons.
o Most abundant element is Argon (found in nature)
d-Block Elements: Groups 3-12
 Considered the transition metals.
 As a group, they are generally good conductors and have a luster.
 They are less reactive than s-block metals.
o Some are so unreactive that they seldom form compounds.
 Major group are the coinage metals.
f-Block Elements: Lanthanides and Actinides
 Are considered the innertransition metals.
 f-block is wedged between group 3 and 4
 The actinides are all radioactive, and every element past Np is artificially made.
Hydrogen and Helium:
 Both elements are light elements that could lift a balloon.
o Remember the Hindenberg?
5-3 Electron configuration and Periodic Properties
Objective: Further explores the relationship between the periodic law and
electron configuration, including trends in properties of electron affinity,
electronegativity, ionization energy, atomic radii, and ionic radii.

Atomic Radii – one-half the distance between the nuclei of identical atoms that
are bonded together.
o Atom size increases going down a group because electrons are being
added to higher energy sublevels. (don’t worry about memorizing
exceptions).
o Atom size decreases going across a periods left to right.
 This is opposite what you would intuitively think.
 Why? As you move across a period, the atom gains electrons as
well as protons. Because the effective nuclear charge has
increased, the electrons will therefore be pulled closer to the
nucleus, causing the atom to become smaller.
 Could you explain C and N sizes?
o See fig. 5-14 pg. 142.

Ionization Energy – energy required to remove one electron from an atom.
o 1st ionization energy, 2nd ionization energy, 3rd ionization energy, …
 A + energy  A+ + eo Ion – an atom or group of bonded atoms that has a positive or negative
charge.
 Ex. Na+, Cl-, O2-, Fe3+, NH4+, CO32-, …
 You have to memorize a lot of these ions…sorry.
 Monatomic ion – ion containing one atom.
 Polyatomic ion – ion containing atoms bonded together.
o Noble Gases have the highest ionization energies because they have a full
octet.
o Group IA have the lowest ionization energies.
 Ionization energy increases going across a period left to right and
decreases down a group
 Decreases down a group because electrons in order energy
levels are easier to remove.
 Increases across period because more electrons in an
energy level make it more difficult to get to the magic
number of 8.
o It is possible to remove electrons from ions.
 Ex. Cu+ + energy  Cu2+ + e Each successive electron removal from an ion feels an increasingly
stronger effective nuclear charge.
 Look at trend on Table 5-3 pg. 145.
 Practice Problems pg. 146.

Electron Affinity – the energy change that occurs when an electron is acquired by
a neutral atom.
o Most atoms release energy when they acquire an electron.
 A + e- → A- + energy
 If energy is released, use negative sign.
 If energy is absorbed, use positive sign
 Common unit is kJ/mol
o Generally, the halogens gain electrons more readly; therefore, their values
are most negative.
 This fact helps to explain their great reactivity.
o Notice electron affinities become more negative going across period.
(Except between group 14 and 15)
 See fig. 5-17
o Generally, electron affinities become more negative going up a group.
 See fig. 5-18
 There are many exceptions to this general rule.
o It is also more difficult to add an electron to an already negative ion, so the
second electron affinities are all positive.
 Ex. Cl- takes less energy to form than Cl2
O2- takes less energy to form than O3-

Ionic Radii – the size of an ion.
o Cation – a positive ion: Na+, Mg2+, Fe2+, Fe3+, …
 Formed by the loss of an electron.
 Cations are smaller than their neutral atom.
 As an electron is lost, the effective nuclear charge is
greater, causing the electron cloud to become pulled closer
to the nucleus, making it smaller.
o Anion – a negative ion: Cl-, F-, N3-, …
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Formed by the gain of an electron(s).
Anions are larger than their neutral atom.
 As an electron is added, the nucleus remains the same so
the electrons are not drawn as close to the nucleus. Also,
added electron causes repulsion among other electrons.
Ion radius increases going left to right across a period and top to
bottom in a group.

Valence Electron – refers to the outer most electrons of an atom.
o These are the electrons that cause chemical reactivity.
o Number of valence electrons an atom has can be determined from
the group number.
 Group IA – 1
+1
 Group IIA – 2
+2
 Group IIIA – 3
+3
 Group IVA – 4
+ or - 4
 Group VA – 5
3 Group VIA – 6
2 Group VIIA – 7
1 Group VIIIA – 8
0 or 8
o See Table 5-4

Electronegativity – a measure of the ability of an atom in a chemical
compound to attract electrons.
o In a chemical compound, the more electronegative atom will
attract electrons more.
 Ex. HCl : H-Cl : electrons are pulled closer to Cl
o The most electronegative atom on the PT is F.
o Electronegativity tends to increase across a period.
o Electronegativity tends to decrease down a group.
 See fig, 5-20 and 5-21 pg. 151-152.
Periodic Properties of the d- and f-block Elements
 The properties of these two blocks are much more difficult to analyze
because of the extreme variability.
o Remember – s-block = 2 electrons, p-block = 6 electrons
d-block = 10 electrons, f-block = 14 electrons
 Atomic Radii
o Generally decrease across a period for d-block, but there is much more
variability because of electron repulsion.
 As the number of electrons increases (period 6 and 7), the size
actually increases because of the electron repulsion.
 See fig. 5-14
o F-block radii is more difficult to predict. I won’t ask you about
specifics.

Ionization Energy
o Generally increase right across a period.
o Generally increase going down a group.
 This is different from representative elements.
 This is because the electrons available for ionization in the
outer s sublevels are less shielded from the increasing nuclear
charge by electrons in incomplete (n - 1)d orbitals.

Ion Formation and Ionic Radii
o The order in which electrons are removed from all atoms of d- and fblock elements is exactly the reverse of the order given by the electron
configuration.
 Ex. Look at electron configuration for Fe. Can you explain
why a Fe3+ ion is possible?

Electronegativity
o Both blocks follow the same rule as representative elements.