Periodic Properties of the Elements

Chapter 7 AP Chemistry

I History and Development

Elements in pure form like Gold and Silver were known thousands of years ago

31 elements were known in 1800 by 1865 the number had more than doubled to 63.

The increase lead to the grouping and classification of the elements

First Periodic Law

Properties of elements are a periodic function of atomic weight

Although Meyer also grouped elements like Mendeleev

Mendeleev was given more credit because he left blanks for un-discovered elements

Mendeleev

Modern periodic law In 1913, by using x-ray

spectra obtained by diffraction in crystals, he found a systematic relation between wavelength and atomic number,

Using Atomic number Mosley rearranged the elements noting that the elements were a periodic function of atomic number

Henry Mosley

Design of the Table

Rows and Columns

Horizontal rows – series or periods – numbered 1-7 Notes the shell or principle QN

Vertical columns – group or a family – numbered 1-18 or 1A-8A, 1B-8B

Review electron configuration notation

Periodic Trends

Size of Atoms Within each group the atomic radius tends to

increase as you go down the group – the principle QN increases – size of the cloud gets larger

Within a period the atomic radius decreases as you go from left to right – the effective nuclear charge increases – the number of protons increases while the number of core electrons stays the same – the outer electrons don’t shied each other well

Ionization Energy (I)

I – the energy required to remove the outermost electron from a gaseous atom or ion

I1 – the energy to remove one electron – I2 – the second electron

Metals tend to give up e- to obtain a Nobel gas configuration

Ionization Energy Trends

Ist Ionization energy generally follows effective nuclear charge – increases from left to right in periods and from bottom to top in groups

Except from group 2A to 3A there is a slight increase – the second e- in the S sub-shell is harder to remove that the 1st P e-

GR 5A to 6A there is a slight increase due to repulsion of paired e-s in the P4 configuration

Every element shows large increase in IE when e’s are removed from a Nobel gas core

Electron Affinities

Is the energy change of the reaction of adding an electron to a gaseous atom or ion.

Tends to be an exothermic processes (Delta H is neg) although in some cases it is positive or endothermic

In general electron affinity tends to decrease ( become more negative) from right to left in periods. There is very little change going down a group in the value.

Properties of Metal, Non-Metals & Metalloids Metals – conduct heat

and electricity are lustrous, malleable, and ductile

Tend to lose e’s to become cations

Form basic oxides – metal oxides react with water to produce basic or alkaline solutions

Ex 7.7,7.8 in text

Metal Properties Cont.

Metal oxides + an acid yield a salt and water

Ex 7.9,7.10 in the text

Non-Metal Properties

Poor conductors of heat and electricity are often dull in color and shatter if forged

Nonmetals tend to gain electrons form other sub and become anions

Non-metals from acidic oxides – non-metal oxide react with water to form acidic solutions

Ex 7-12, 7-13 in text Non-metal oxides + a base yield a salt plus

water ex 7.14, 715 in text.

Mettaloids

Boundary between metals and non-metals Mettaloids can either gain or lose electrons Al and Po are metals not mettaloids Properties between metals and non-metals Several are semiconductors and re important

because of their use in circuits and computer chips

Metallic character – increases from left to right and from top to bottom in a group

Group Trends

Alkali Metals Soft, gray metals Low Ionization energies

– lose e’s easily – form +1 cations – very reactive

Electrolysis is a technique to force an electron back on a cation to produce a neutral metal – the metals can be obtained

Alkali Metals

by passing an electrical current through a molten salt. The metals combine with hydrogen to from

hydrides ex. 7.19, or with Sulfur to form Sulfides ex. 7.20 or with chlorine to form chlorides. ex 7.21

The alkali metals react with water to produce hydrogen gas and hydroxides ex 7.22.

See examples of reactions with oxygen ex 7.23-7.25

Alkaline Earth Metals

Slightly harder and more dense than Alkali metals

1st ionization energies are low but not a s low as alkali metals form 2+ cations

Reactivity tends to increase as you progress down a group.

See examples of reactions 7.26 – 7-29

Got Milk?

Trends in Selected Non-metals

Hydrogen Usually placed in grp 1A because of its 1s

electron config. Although a non-metal It reacts with other non-metals to form

molecular cmps ex 7.30-32 Reacts with metals to form hydrides 7.31-32

Oxygen Family

Non-metallic – top of family – metallic at bottom

O, S, Se are typical non-metals – O and S are allotropes – molecules of elements with different structures ( O2, O3, S2, S4, S6, S8)

Polonium is a rare and radioactive element See examples of reactions 7.33-36

The Halogens

All non-metals (astatine is radioactive – seldom discussed)

All are diatomic in gas phase

Bromine normally a liquid – iodine is normally a solid

See reactions 7.37-42

Nobel Gases

Mostly non-reactive Some compounds of

Xenon and Krypton are known

All monatomic gases Highest ionization

energies of any family

xenon tetrafluoride