ch.1 Periodic Structure and Trends

Horizontal rows on the periodic table

Periods

Vertical columns on the periodic table

Groups or families

Elements with similar chemical properties

Elements in the same group

Electrons that mainly determine group behavior

Valence electrons

Two separated rows at the bottom

Lanthanide and actinide series

Why do elements in the same group behave similarly?

They have similar valence-electron configurations.

How do periods and groups differ?

Periods are horizontal rows; groups are vertical columns.

Far-left and center elements that conduct well

Metals

Far-right elements that conduct poorly

Nonmetals

Elements with both metallic and nonmetallic traits

Metalloids

Another name for metalloids

Semimetals

Common physical feature of metals

Shiny

Electrical behavior of metals

Good conductors

Electrical behavior of nonmetals

Poor conductors

How do metals, nonmetals, and metalloids differ?

Metals conduct well and usually lose electrons; nonmetals conduct poorly and often gain or share electrons; metalloids have properties of both.

Group 1 elements except hydrogen

Alkali metals

Usual ion formed by Group 1

+1 cation

Reason Group 1 forms +1 ions

Loss of one valence electron

Group 2 elements

Alkaline earth metals

Usual ion formed by Group 2

+2 cation

Reason Group 2 forms +2 ions

Loss of two valence electrons

Groups 3 through 12

Transition metals

Group known for multiple oxidation states

Transition metals

Why are Roman numerals often used for transition metals?

They can form ions with different charges.

How do alkali metals, alkaline earth metals, and transition metals differ?

Alkali metals usually form +1 ions, alkaline earth metals form +2 ions, and transition metals can have multiple oxidation states.

Group containing boron

Group 13

Usual charge associated with Group 13

+3

Group 14 name

Carbon family

Valence electrons in Group 14

4

Group 15 name

Nitrogen family

Valence electrons in Group 15

5

Group 16 name

Chalcogens

Usual ion formed by Group 16

−2

Group 17 name

Halogens

Usual ion formed by Group 17

−1

Group 18 name

Noble gases

Another name for Group 18

Inert gases

Reason noble gases are mostly unreactive

Filled valence shell

Why are halogens highly reactive?

They need only one electron to fill their valence shell.

How do Groups 16, 17, and 18 differ?

Group 16 often gains two electrons, Group 17 gains one, and Group 18 already has a filled shell.

Attractive pull of the nucleus felt by valence electrons

Effective nuclear charge

Symbol for effective nuclear charge

Zeff

Cause of nuclear attraction

Positive protons attracting negative electrons

Trend of Zeff across a period

Increases from left to right

Reason Zeff increases across a period

More protons are added without much additional shielding.

Trend of Zeff down a group

Decreases in the simplified trend from the notes

Reason valence electrons feel less nuclear pull down a group

More shells increase distance and shielding.

Blocking of nuclear attraction by inner electrons

Electron shielding

How do effective nuclear charge and shielding differ?

Effective nuclear charge is the pull felt by valence electrons; shielding reduces that pull.

Size of a neutral atom

Atomic radius

Atomic-radius trend down a group

Increases

Atomic-radius trend across a period

Decreases from left to right

Reason atomic radius increases down a group

Additional electron shells and greater shielding

Reason atomic radius decreases across a period

Stronger effective nuclear charge pulls electrons inward.

Direction of increasing atomic radius

Down and left

How does atomic radius relate to effective nuclear charge?

A stronger effective nuclear charge pulls electrons closer and decreases atomic radius.

Size of an ion

Ionic radius

Effect of gaining electrons on radius

Increases radius

Effect of losing electrons on radius

Decreases radius

Ion larger than its neutral atom

Anion

Ion smaller than its neutral atom

Cation

Why is an anion larger than its neutral atom?

Extra electrons increase electron-electron repulsion.

Why is a cation smaller than its neutral atom?

Fewer electrons reduce repulsion, and an outer shell may be lost.

How do cation, neutral-atom, and anion sizes compare?

Cations are generally smaller than neutral atoms, while anions are generally larger.

Energy required to remove an electron

Ionization energy

Ionization-energy trend across a period

Increases from left to right

Ionization-energy trend down a group

Decreases

Direction of increasing ionization energy

Up and right

Why does ionization energy increase across a period?

Stronger effective nuclear charge holds electrons more tightly.

Why does ionization energy decrease down a group?

Valence electrons are farther away and more shielded.

How do atomic radius and ionization energy usually relate?

Smaller atoms usually have higher ionization energy because their electrons are held more tightly.

Energy needed to remove the first electron

First ionization energy

Energy needed to remove a second electron

Second ionization energy

Which is higher: first or second ionization energy?

Second ionization energy

Why is the second ionization energy higher?

After one electron is removed, the remaining electrons experience less repulsion and stronger attraction.

Why is sodium’s first ionization energy relatively low?

Removing one valence electron gives sodium a stable noble-gas configuration.

Why is sodium’s second ionization energy extremely high?

It would remove an electron from a stable filled inner shell.

How do first and second ionization energies differ?

The first removes an electron from the neutral atom; the second removes one from an already positive ion and is harder.

Character that increases down and left

Metallic character

Character that increases up and right

Nonmetallic character

Trend shared by ionization energy and nonmetallic character

They generally increase up and to the right.

Trend shared by atomic radius and metallic character

They generally increase down and to the left.

An element has a large radius and low ionization energy. Where is it likely located?

Near the lower-left portion of the periodic table.

An element has a small radius and high ionization energy. Where is it likely located?

Near the upper-right portion of the periodic table.

Which is larger: Na or Na+?

Na

Which is larger: Cl or Cl−?

Cl−

Which has a higher ionization energy: Na or Cl?

Cl

Which has a larger atomic radius: Na or Cl?

Na

Why is Cl smaller than Na even though both are in the same period?

Chlorine has more protons and a stronger effective nuclear charge.

Why is potassium larger than sodium?

Potassium has an additional electron shell.

Why is fluorine harder to ionize than lithium?

Fluorine has a stronger effective nuclear charge and a smaller atomic radius.