General Chemistry I

Chapter 7

Periodic Properties of the Elements

Development of the Periodic Table

n    Most elements exist in compounds

n    Difficult to isolate

n    19th Century saw huge increase in the number of elements discovered

n   From 31 in 1800 to 63 in 1865

n    Mendeleev and Meyer

n    Arrange based on properties and in increasing atomic weight

n    Mendeleev considered ‘Father of the Periodic Table’

n    Was more adamant about it’s use

n    Left ‘holes’ for missing elements

n   Ga and Ge

n    Used his table to predict the properties

n    When elements were discovered, predicted properties were very close to observed.

n    Problems with atomic weight arrangement

n    K < Ar and I < Te

n    Mendeleev stated when techniques became more accurate, these problems would resolve themselves

n    Moseley

n    Studied X-ray emission

n    Moseley found that emitted X-rays were of different frequencies for different elements.

n    Generally frequency increased as the atomic mass increases

n    Exceptions were Ar, K and Te, I

n    First to assign an ‘Atomic number’

n   Related to number of protons

n   Helped to identify holes

Electron Shells and Atomic Size

n    Atom does not have ‘sharp’ boundaries

n    Often define a ‘probability surface’ (typically 90%) that is a hard surface

n    Electron Shells

n    First proposed by Lewis

n   Similar to an onion

n   Predates Bohr

n    Analogy in QM atom

n     Examine noble gas electronic structure

n    He – 1s2

n    Ne – 1s2 2s2 2p6

n      Ar – 1s2 2s2 2p6 3s2 3p6

n     What is the ‘radial electron density’ for these elements?

Electron Density

n    Supports idea of shells

n    Considerable overlap of the shells

n    Not as ‘clean’ as an onion

n    Why is first shell closer to the nucleus for Ar than Ne

n    Higher nuclear charge

n    Electrons held ‘tighter’

Atomic Size

n     Non-bonding radius

n    Closest atoms will approach before electron-electron repulsion.

n     Bonding atomic radius

n    Distance between two like nuclei in a chemical bond.

n    Many atomic properties depend on atomic size.

n    More later…

n    Trends

n    Within a group, size increases as we go down a group.

n    Within a period, size decreases as we go across (left to right) a period.

n    Two factors

n    Principal QN

n   Explains why size increases top to bottom

n    Effective Nuclear Charge

n   Nuclear charge increases L-R

n   Screening remains constant

n   Electrons held tighter (closer)

Problems:  1 – 15 odd

Ionization Energy

n    Energy required to remove an electron from an atom/cation

n    First electron removed, 1st ionization energy

n    Second electron, 2nd ionization energy

n    etc…

n    Large jump occurs when you start removing core electrons.

First Ionization Energies

n     1st ionization energies decrease from top to bottom in a family.

n     1st ionization energies generally increase from L-R across a period.

n     Anomolies

Electron Affinities

n    Energy change that occurs when an electron is added to an atom in the gaseous state.

n    The more negative the electron affinity, the greater the attraction for an electron

n   Negative value means energy is released.

n   Energy may be released or absorbed when an electron is added.

Periodic Trends

n     Not as ‘clean’ as ionization energies

n     Generally become less negative from top to bottom

n     Left to right generally more negative

n    MANY exceptions

Problems:  17 – 29 odd

Metals vs. Nonmetals

n    Properties generally opposite

n    Metallic character

n    Nonmetallic character


n     Luster

n     Conductors

n     Malleable (sheets)

n     Ductile (wires)

n     Solids, except Hg

n    MP from –39OC to 1900OC (Cr)

n     Low ionization energies

n    Form cations

n    Transition metal cations often have variable charge

n    Many have +2 charge

n    Loss of ns electrons

n     Metal/nonmetal generally form ionic compounds

n     Soluble metal oxides basic

n    Na2O+H2O ->2NaOH

n    Basic character of oxides also demonstrated by reaction with acids


n    Vary greatly in appearance

n     Non lustrous

n     Poor conductors

n     MP generally lower

n    Diamond @ 3570OC

n     May be s, l, or g

n    Solids range from very soft to very hard

n     Tend to form anions

n     Form molecules with other nonmetals

n     Nonmetal oxides are acidic

n    CO2+H2O -> H2CO3

n    CO2, SO2, SO3, NOx main contributors to acid rain


n    Some metallic/some nonmetallic characteristics.

Problems:  31 – 39 odd

Group Trends – Active Metals

n    Alkali Metals

n    Group IA

n    First isolated from ashes

n    Always in nature as compounds or ions

n   Na+ / K+ important for intracellular transport

n    Chemistry dominated by loss of single electron

n   M -> M+ + e-

Alkali Metals

n    Pure metals formed by electrolysis

n    Na from molten NaCl (801OC)

n    Form hydrides with H

n    Donates electron to hydrogen to form H-

n    Combine directly with nonmetals to form ionic compounds

n    React with water to produce H2(g)

n    Very exothermic

Reaction with Oxygen

n    Li – forms Li2O

n    Reacts in water to form lithium hydroxide

n    Other alkali metals form peroxides

n    Na-O-O-Na (sodium peroxide)

n   The peroxide bond is quite unstable!

n    Salts not colored unless combined with a colored anion.

n    Will emit characteristic colors in a flame.

Alkaline Earth Metals

n    Similar to alkali metals

n    Harder

n    Higher density

n    Higher MP

n    Chemistry dominated by the loss of 2 electrons

n    Not as reactive as alkali metals

n    Metals are more stable when refined

n    Mg protected by coating of MgO

n    Can be used as building material

n    ‘Mag’ wheels

n    Heavier IIA metals must be protected from air/water.

n    Mg and Ca essential for living organisms.

Group Trends - Nonmetals

n    Hydrogen

n    Should be considered it’s own family!

n    Unique because both loss of 1 electron and gain of 1 electron gives a stable configuration.

n   I1 typical of nonmetals

n    Chemistry dominated by H+ ion

n    Chem II

Group VI A

n    Chalcogens

n    Oxygen is a gas – all others solids

n    From nonmetal to metal

n    Molecular oxygen – 2 forms

n    ‘oxygen’ (O2) and ozone (O3)

n   Allotropes

n    Will ‘oxidize’ most other elements

n    Oxygen

n    Forms oxides (O2-), peroxides (O22-) and superoxides (O2-)

n   Often decompose to form O2

n    Sulfur

n    Yellow

n    Most common form S8

n    Tends to gain electrons to form sulfides

n    Chemistry more complex than oxygen


n    ‘Salt formers’

n    All nonmetals

n    All diatomic molecules

n    Range from solid (F2, Cl2) to liquid (Br2) to solid (I2)

n    All colored compounds

n    Chemistry dominated by the formation of anions by gaining electrons.

n    Fluorine

n    Will remove electrons from anything

n    Very exothermic

n   Used for uranium enrichment and glass etching

n    Chlorine

n    Most used industrially

n    Other product of electrolysis of molten salt

n    Reacts slowly with water

n   Forms HCl and HOCl

n   Used as disinfectant for water

n    Form soluble halides with most metals

n    Bromine

n    One of two liquid elements at room temp.

n    Reactions similar to chlorine

Noble Gases

n    Monoatomic

n    Characterized by stability

n    Valence shells s and p orbitals filled

n    Very few compounds of noble gases known

n    Those that are known are for the heavier noble gases (lower I1) with fluorine.

Problems:  41 – 55 odd