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

Put the elements into groups according to properties

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

Germanium

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

Mendeleev’s predictions

Actual element

Atomic mass = 72 Atomic mass = 72.60

High melting point Melting point = 958 °C

Density = 5.5 g/cm3 Density = 5.36 g/cm3

Dark gray metal Gray metal

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?

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?

Nickel and cobalt Argon and potassium Better mass measurements just

confirmed the discrepancy

Explanation

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

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

Not discovered on Earth until 1894 - 1900.

Group 18 was added to the table

Lanthanides

Hard to separate All have similar properties Added to the table in the early

1900s

Actinides

Discovered later Also all have similar properties

Periodicity

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

Chemical properties of an element

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

Period length

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

Table 5-1

Full periodic table

Table with f-block in place

1st period

1s sublevel being filled 1s can hold 2 electrons, so there are

2 elements in the 1st period.

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

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

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

Shows blocks

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

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

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

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 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

Discuss

Page 133 Sample problem 5-1 and practice

problems

Discuss

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

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

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

End in p1 to p6. Coefficients are the same as the period

ns2np1

Always have a full s-sublevel

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

Halogens

Group 17 Most reactive nonmetals Form compounds called salts

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

Discuss

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

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

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

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.

Discuss

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

Ion

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

Ionization

Any process that makes ions

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

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

Energy required to remove other electrons from positive ions.

IE2, IE3, etc Get higher as you remove more

electrons Less shielding

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

Discuss

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

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

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 trends

Not as regular Usually, electrons add with greater

difficulty as we move down

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.

Discuss

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

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

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 trends

Ionic radius increases as you go down the table

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

Electronegativity

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)

Increases left to right across the rows

Decreases down the columns

Discuss

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

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 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

Discuss

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

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.