Perodic Trends

8/14/2019 Perodic Trends

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Trends of the

Periodic Table


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Review! Periodic Table was first organized by

Dmitri Mendeleev in the mid 1800s

Mendeleev organized the elements by chemicalreaction in rows, then by atomic mass in columns

Henry Moseley then took Mendeleevs table, keptthe chemical reactivities together, but placed themin columns instead. He also ordered the elements

by increasing atomic number in rows.

When Moseley did this, all the periodic trends justfell into place.

Remember: columns = groups/families,rows = periods


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Periodic Trends


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Electrons

Electrons do not freely float in space

Orbit around nucleus: Electron shells

Each shell corresponds to an amount of

energy.


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Valence

Electrons The valence electrons are the outermost electrons of an atom. The valence electrons determine the chemical properties

Number of valence electrons equals the column number in theA columns

Elements with the same number of valence electrons are very

similar chemically

Alkali metals in Group 1A 1 valence electron

Li, Na, K, Rb, Cs

Halogens in Group 7A 7 valence electrons

F, Cl, Br, I


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Atomic Radius What is Atomic Radii?

Distance from thenucleus to theoutermost level of e-(aka the valence shell)

What trend do you seeas you go across (leftto right) the period?

Atomic radius

decreases Down the group?

Atomic Radiusincreases

WHY???


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Explaining the Trend

As you go L to R, the atomic radius decreases

because as you go L to R, the amount of attraction

between p+

and e-

increase.More attractions = smaller atomic radius

As you go down a column, atomic radius increases

because the e- are farther away from the nucleus.

There are weaker attractions.

Weaker attractions = larger atomic radius


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Electronegativity

What is Electro-negativity?

An atoms Luuuvvvfor electrons!

The tendency toattract another

atoms electrons What trend do you see

as you go across theperiod?

Electronegativity

increases! Down the group?

Electronegativitydecreases!

WHY???


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As you go L to R, electronegativity increasesbecause of the increase in protons. The moreprotons, the more able it will be to attract otheratoms electrons.

More attractions (small radius) = large electronegativity

As you move down a column, electronegativitydecreases because of the increase in numberelectron an atoms already has. This means theatom will be less able to attract another atomselectrons.

Less attractions (large radius) = small electronegativity


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Ionization Energy What is Ionization

Energy? The energy needed to

remove an electron

What trend do you see

as you go across theperiod?

Ionization E increases

Down the Group?

Ionization E decreases

WHY???


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As you go L to R, the ionization energy increases becauseof the increase in the number of protons. The more

protons, the more energy that is needed to remove an

electron.More attractions (small radius) = large ionization energy

As you go down a column, the ionization energy decreasesbecause of the decrease in attractions.

Due to electron shielding More electrons, leads to outer electrons less tightly

held.

The less attractions, the lower the energy that is needed toremove an electron.

Less attractions (large radius) = small ionization energy


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Ionization Energy

Amount of energy required to remove an

electron from the ground state of a gaseous

atom or ion.First ionization energy is that energy required

to remove first electron.

Second ionization energy is that energyrequired to remove second electron, etc.


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Ionization Energy

It requires more energy to remove each successive

electron.

When all valence electrons have been removed, theionization energy takes a quantum leap.


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Electron Affinity What is Electron

Affinity? The energy needed to

add an electron

As you go across the

period electron affinityincreases .

Electron affinitydecreases down the

family WHY???


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Explaining the trend

As you go L to R, the electron affinity increases because ofthe increase in the number of protons. The more protons,the greater the attraction the protons have for electrons.

More attractions (small radius) = large electron affinity As you go down a family, the electron affinity decreases

because of the decrease in attractions.

Due to electron shielding

More electrons, leads to outer electrons less tightlyheld.

The less attractions, the lower the electron affinity

Less attractions (large radius) = small electron affinity


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