Periodic Properties of the Elements

Chapter 7

The Periodic Table

Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s).

Found that similar chemical and physical properties recur periodically when the elements are arranged in order of increasing atomic number.

At this time they didn’t know about atomic numbers, so used masses which generally increases as atomic # increases.

Had blanks in the table and used periodicity to guess about the characteristics of the missing elements.

Moseley

After Rutherford proposed the nuclear model of the atom, Henry Moseley developed the concept of atomic numbers.

Bonding Atomic Radius- based on the distance separating atoms when they are chemically bonded to one another

This radius is shorter than the non bonded radius due to the nuclear attraction between the two atoms.

The bonding atomic radius decreases as you go across the period, and increases as you go down a group. (Rb> F)

Bond length

Predict which will be greater, the P-Br bond length in PBr3 or the As-Cl bond length in AsCl3

P-Br

Radial Electron Density

The probability of finding an electron with respect to the nucleus

The 1s subshell in Ar is much closer to the nucleus than the 1s subshell of He. This is because of the Zeff

Ionic Radius

Cations are formed when metallic atoms lose valence electrons.

These ions have smaller radii than their parent atoms

Anions are formed when nonmetallic atoms gain electrons

These ions are larger than their parent atoms due to the extra repulsions of another electron

Periodic Trends

How easily an electron will be removed from an atom is an important indicator of the chemical nature of that atom.

Ionization energy is the energy required to remove an electron from the ground state a gaseous atom

The greater the ionization energy, the more difficult it is to remove an electron

Ionization Energy

Highest energy electron removed first (outermost).

First ionization energy (I1) is that

required to remove the first electron. Second ionization energy (I2) - the

second electron etc. etc.

Trends in ionization energy

for Mg • I1 = 735 kJ/mole• I2 = 1445 kJ/mole• I3 = 7730 kJ/mole

The effective nuclear charge increases as you remove electrons.

It takes much more energy to remove a core electron than a valence electron because there is less shielding.

This trend is because the positive nuclear charge that provides attractive forces remains the same, while the number of electrons which provide repulsive forces decreases.

Explain this trend

For Al• I1 = 580 kJ/mole

• I2 = 1815 kJ/mole

• I3 = 2740 kJ/mole

• I4 = 11,600 kJ/mole

Ionization Trends

Generally from left to right, I1 increases

because there is a greater nuclear charge with the same shielding. (Generally, the alkali metals show the lowest ionization energies in a row, and the noble gases the highest.

As you go down a group I1 decreases

because electrons are farther away.

It is not that simple

Zeff changes as you go across a period,

so will I1

Half filled and filled orbitals are harder to remove electrons from.

here’s what it looks like.

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Atomic number

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Electron Affinities

The energy change that occurs when an electron is added to a gaseous atom to form a negative ion. (A measure of the affinity or attraction for the added electron.

For most atoms the electron affinity is negative because energy is released when an electron is added.

Different entities

Remember!!! Ionization Energy is the desire to lose an

electron (+) Electron Affinity is the desire to gain an

electron

Electron Affinity Trends

Generally becomes increasingly negative as you go toward the halogens. For the noble gases, the electron affinity is positive, meaning the ion will not form because that would mean that the gas would have to go to a higher energy sub shell which is energetically unfavorable.

Any time the value is zero, the ion will not form. The bigger the negative, the more likely that the ion will form.

Chlorine has the highest electron affinity.

Parts of the Periodic Table

Metals, Non metals, and Metalloids

Metals

Roughly 3/4 of the elements are metals. Properties of metals include luster,

malleability, ductility, god conductors of heat & electricity, form cations in an aqueous solution.

The more an element shows properties of metals, the greater it’s metallic character

(Increases down a row, decreases across a period)

More Properties of Metals

All metals are solid except for mercury Metals tend to have low ionization

energies which is why they are oxidized (lose electrons) to form a cation when they undergo a chemical reaction.

When the outermost electrons are lost, the ion achieves a noble gas configuration

Many of the transition metals have the ability to form more than one ion

Metal Oxides

Metal-nonmetal compounds are said to be ionic

Oxides are especially important (oxygen is everywhere!)

Most metal oxides are basic, which means they dissolve in water to form bases

Metal oxides can also react with acids to form salt + water

Nonmetals

Poor conductors of heat & electricity Vary in appearance Have lower melting points than metals Several exist as diatomic molecules Tend to gain electrons in a chemical

reaction to fill their outer p subshell completely giving a noble gas configuration

More properties of Nonmetals

Nonmetal bonded to a nonmetal makes a molecular substance

These molecules tend to be gases, liquids, or low-melting solids.

Non-metal oxides are generally acidic which means they combine with water to form an acid. (This is why carbonated water is acidic)

These acidic nonmetal oxides will combine with a base to produce salt & water

Metalloids

Have properties of both metal and nonmetal!

Silicon- looks like a metal, brittle like a nonmetal, semiconductor used in computer chips

Group Trends

Active Metals

The Alkali Metals

Group 1 Soft metallic solids Doesn’t include hydrogen- it behaves as

a non-metal Down the group-decrease in IE Down the group-increase in radius Decrease in density Decrease in melting point

Alkali Metals

For each row, the alkali has the lowest ionization energy

All very reactive and lose 1 electron to form +1 cation

Exist in nature only as compounds Electrolysis used React vigorously with water to produce

hydrogen gas and metal hydroxides Exothermic enough to ignite Hydrogen

Also extremely reactive with oxygen Stored in kerosene Do not produce a colored solution

because no electron to excite

Alkaline Earth Metals

All solids Harder , more dense, melt at higher temps

than alkali Slightly higher ionization energies, thus

slightly less reactive (compared to Alkali). Increasing reactivity as you go down the group

that accounts for why Berylium does not react with water but Calcium and everything below it do.

Tend to lose 2 electrons and form +2 cation

Alkaline Earth Metals

Highly reactive so usually found in nature as part of a compound

Group Trends

Nonmetals

Hydrogen

Nonmetal that occurs as a colorless diatomic gas

Since there is no shielding, has an extremely high ionization energy

Usually combines with other non-metals to form molecular compounds

Reacts with active metals to form metal hydrides (H is -)

Oxygen’s Group

Changing trends as you go down the group

Oxygen usually found in two molecular forms oxygen and ozone (allotropes- different forms of the same element in the same state)

Oxygen makes up 21% of air Ozone is toxic and smelly Oxygen usually present as the oxide ion

Oxygen’s Group

Sulfur (exists as eight membered rings of sulfur atoms

Most sulfur is found as metal sulfides Can be burned in oxygen to produce

sulfur dioxide (pollutant)

The Halogen Family

(Astatine omitted because extremely rare, radioactive and unknown)

All typical non-metals Melting and boiling point increase as you

go down the group All diatomic Tend to gain electrons and form -1 anion Have highly negative electron affinities Fluorine and Chlorine most reactive

The Noble Gases

All non-metal gases All monoatomic Rn too highly radioactive to study Completely filled s and p subshells Large ionization energies which

decrease as you go down the group Inert gases because thought to be

unable to form compounds