The periodic law
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The periodic law. Chapter 5. Why do we need a table?. To organize the elements To show trends. Periodic. A repeating pattern. Mendeleev’s table. 1869 – Dmitri Mendeleev – Russian Arranged the elements in order of increasing mass and noticed that chemical properties were periodic

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The periodic law

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The periodic law

The periodic law

Chapter 5

Why do we need a table

Why do we need a table?

  • To organize the elements

  • To show trends



  • A repeating pattern

Mendeleev s table

Mendeleev’s table

  • 1869 – Dmitri Mendeleev – Russian

  • Arranged the elements in order of increasing mass and noticed that chemical properties were periodic

  • Put the elements into groups according to properties

Mendeleev vs meyer

Mendeleev vs. Meyer

  • 1860s Mendeleev and German Lothar Meyer each made an eight column table.

  • Mendeleev left some blanks in his table in order for all the columns to have similar properties – he predicted elements that hadn’t been discovered yet.

Why similar properties

Why similar properties?

  • Why did they group according to properties and mass and not atomic number or number of outer level electrons?



  • Mendeleev’s blank spots and his ability to predict future elements helped his table win acceptance.

Mendeleev s table1

Mendeleev’s table

  • Elements arranged in order of increasing mass.

  • Properties are repeated in an orderly, periodic, fashion.

  • Mendeleev’s periodic law – the properties of the elements are a periodic function of their masses.

Mass mistakes

Mass mistakes?

  • In order for Mendeleev to arrange his elements by properties, he had to put tellurium and iodine in the wrong order.

  • He explained this by assuming that their masses hadn’t been measured very accurately.

More mass mistakes

More mass mistakes?

  • Nickel and cobalt

  • Argon and potassium

  • Better mass measurements just confirmed the discrepancy



  • 1913 – Henry Moseley

  • X-ray experiments revealed the atomic number was the number of protons

  • Modern periodic law – the properties of the elements are a periodic function of their atomic numbers

Modern periodic table

Modern periodic table

  • An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.

Noble gases

Noble gases

  • Not discovered on Earth until 1894 - 1900.

  • Group 18 was added to the table



  • Hard to separate

  • All have similar properties

  • Added to the table in the early 1900s



  • Discovered later

  • Also all have similar properties



  • Elements in the same group (column) have similar properties.

Chemical properties of an element

Chemical properties of an element

  • Are governed by the electron configuration of an atom’s highest energy level

Period length

Period length

  • Determined by the number of electrons than can occupy the sublevels being filled in that period.

  • Table 5-1

Full periodic table

Full periodic table

  • Table with f-block in place

1 st period

1st period

  • 1s sublevel being filled

  • 1s can hold 2 electrons, so there are 2 elements in the 1st period.

2 nd and 3 rd periods

2nd and 3rd periods

  • 2s and 2p or 3s and 3p being filled

  • s and p sublevels can hold 8 total, so there are eight elements in these periods

4 th and 5 th periods

4th and 5th periods

  • Add d sublevels, which can hold 10 electrons

  • Need to fill 4s, 3d, and 4p – 18 electrons

  • 18 elements in each period

6 th and 7 th periods

6th and 7th periods

  • Add f-block, which holds 14 electrons

  • Fill 6s, 5d, 4f, 6p

  • Need 32 electrons

  • 32 elements in each period

Figure 5 5

Figure 5-5

  • Shows blocks

Electron configurations

Electron configurations

  • Elements in columns 1, 2, and 13-18 have their last electron added in an s or p orbital.

  • Elements in columns 3-12 have their last electron added in a d level.

The s block elements groups 1 and 2

The s-block elements: Groups 1 and 2

  • Chemically reactive metals

  • Group 1

    • Have 1 electron in outer s orbital

      • Coefficient represents period

        • Row 2: 2s1, Row 3: 3s1, etc. (ns1)

  • Group 2

    • Have 2 electrons in outer s orbital

      • Coefficient represents period

        • Row 2: 2s2, Row 3: 3s2, etc. (ns2)

Alkali metals

Alkali metals

  • Metals in group 1

  • Have silvery appearance

  • Soft enough to cut with a knife

  • Not found alone in nature

  • React violently with nonmetals

  • Melting point decreases as you go down the table

Alkaline earth metals

Alkaline-earth metals

  • Group 2

  • Harder, denser, and stronger than alkali metals

  • Higher melting points than alkalis

  • Less reactive

  • Not found alone in nature

Hydrogen and helium

Hydrogen and helium

  • Hydrogen

    • Located above group 1 because of its electron configuration

    • Not really in group 1, because its properties don’t match

  • Helium

    • Has an electron configuration like group 2 elements

    • In group 18 because it is unreactive



  • Page 133

  • Sample problem 5-1 and practice problems



  • Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [Rn] 7s1 is located.

    • Group 1, 7th period, s block

  • Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [He]2s2 is located.

    • Group 2, second period, s block

D block elements groups 3 12

d-block elements: Groups 3-12

  • End in d1 to d10.

    • Coefficients are one less than the period

      • Example: Fe is in the 6th column of transition elements in the 4th period, ends in 3d6

Transition elements

Transition elements

  • Groups 3-12

  • Typical metallic properties

    • Good conductors

    • High luster

  • Less reactive than alkalis and alkaline-earths

  • Some are unreactive enough to appear in nature

P block elements groups 13 18

p-block elements: groups 13-18

  • End in p1 to p6.

    • Coefficients are the same as the period

      • ns2np1

      • Always have a full s-sublevel

P block elements

p-block elements

  • Properties vary greatly

  • Includes all nonmetals except hydrogen and helium

    • Solids, liquids and gases

  • Includes all the metalloids

    • Between metals and nonmetals

    • Brittle solids

    • Semiconductors – can conduct under certain conditions

  • Includes some metals

    • Less reactive than alkalis and alkaline-earths



  • Group 17

  • Most reactive nonmetals

  • Form compounds called salts

F block elements

f-block elements

  • Lanthanides and actinides

    • Endings are f1 to f14

    • Coefficients are two less than the period

  • All actinides are radioactive

  • Those after neptunium are synthetic



  • Sample problems and practice problems on pages 136, 138, and 139

  • With your group first, then join with another group.

  • Do you have any questions?

Atomic radius

Atomic radius

  • Ideally, the distance from the center of the atom to the edge of it’s orbital.

    • But, atoms are “fuzzy”, not clearly defined.

  • Defined as one-half the distance between the nuclei of identical atoms that are bonded together.

Period trends see figure 5 13

Period trends – see figure 5-13

  • As we move from left to right across the table, we gain protons.

  • There is a greater positive charge on the nucleus.

  • This greater charge pulls harder on the outer electrons, pulling them in closer.

  • The atom gets smaller.

Group trends

Group trends

  • As we move down the table, the principle quantum number increases.

  • When the principle quantum number increases, the electron cloud gets bigger.

  • The size of the atoms gets bigger.



  • Which of the elements Li, Rb, K, and Na has the smallest atomic radius? Why?

    • Li, it is highest on the table

  • Which of the elements Zr, Rb, Mo, and Ru has the largest atomic radius? Why?

    • Rb, it is farthest to the left on the table

The periodic law


  • An atom or group of bonded atoms that has a positive or negative charge



  • Any process that makes ions

Ionization energy ie

Ionization energy (IE)

  • First ionization energy (IE1) – the energy required to remove the most loosely held electron.

  • Measured in kJ/mol

Ionization energy see figure 5 15

Ionization energy – see figure 5-15

  • Experimentally determined.

    • From isolated atoms in the gas phase

  • Tends to increase as you move across a row from left to right

    • Why group 1 is most reactive

    • Caused by higher charge

  • Tends to decrease as you move down a column

    • Electrons are farther from nucleus

    • Shielding from inner electrons

Other ionization energies see table 5 3

Other Ionization Energies – see Table 5-3

  • Energy required to remove other electrons from positive ions.

  • IE2, IE3, etc

  • Get higher as you remove more electrons

    • Less shielding

Noble gases1

Noble Gases

  • Have High ionization energies

  • When a positive ion of another element reaches a noble gas configuration, its ionization energy goes up.

    • Example: When K loses one electron, it has Ar’s electron configuration

    • This makes it stable

    • Its IE2 is much higher than its IE1



  • State in words the general trends in ionization energies down a group and across a period of the periodic table.

Electron affinity

Electron affinity

  • The energy change that occurs when an electron is gained by a neutral atom

    • Most atoms release energy

      • Represented by a negative number

    • Some atoms gain energy

      • Represented by a positive number

      • These ions will be unstable

  • KJ/mol

Period trends see figure 5 17

Period trends – see figure 5-17

  • Group 17 has most negative electron affinity.

  • Tends to get more negative (release more energy) as we move to the right

  • Exceptions:

    • groups with full or half-full sublevels are more stable

Group trends1

Group trends

  • Not as regular

  • Usually, electrons add with greater difficulty as we move down

Adding additional electrons

Adding additional electrons

  • Second electron affinities are all positive because it is more difficult to add electrons to a negative ion.

  • If a noble gas configuration has been reached, it is even more difficult.



  • State in words the general trends in electron affinities down a group and across a period of the periodic table.

Ionic radii

Ionic Radii

  • Cation – a positive ion

    • Ionic radius smaller than atomic radius

  • Anion – a negative ion

    • Ionic radius is larger

Period trends see figure 5 19

Period Trends – see figure 5-19

  • Metals form cations by losing electrons

    • Ions are smaller

    • Radius decreases as we move across

  • Nonmetals form anions by gaining electrons

    • Ions are larger

    • Radius decreases as we move across

Group trends2

Group trends

  • Ionic radius increases as you go down the table

Valence electrons

Valence electrons

  • Available to be lost, gained or shared in the formation of chemical compounds

  • In highest energy levels

  • For s-block, the group number is the same as the number of valence electrons

  • For the p-block, the group number is 10 more than the number of valence electrons



  • The measure of the ability of an atom in a compound to attract electrons

    • The atom with higher electronegativity pulls the electrons closer to itself

Electronegativity trends figure 5 20

Electronegativity trends (figure 5-20)

  • Increases left to right across the rows

  • Decreases down the columns



  • Explain why elements with high (more negative) electron affinities are also the most electronegative.

D and f block elements

d- and f-block elements

  • Properties vary less and with less regularity than others

  • Atomic radii

    • d-block

      • Usual patterns

    • f-block (unusual)

      • Increase across periods

      • Decrease down groups

D and f block elements1

d- and f-block elements

  • Ionization energy

    • Increase across periods

    • d-block increases down groups (unusual)

    • f-block decreases down groups

  • Ionic radii

    • Cations have smaller radii

  • Electronegativity

    • d-block follows normal rules

    • f-block all have similar electronegativities



  • Among the main-group elements, what is the relationship between group number and the number of valence electrons?

  • In general, how do the periodic properties of the d-block elements compare with those of the main-group elements?

Prelab notes

Prelab notes

  • Precipitate – solid that falls out of a solution

  • The formation of a precipitate indicates there has been a chemical change.

  • This means that there were ions present that were free to react.

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