Ch. 14: Chemical Periodicity

Standard: Matter consists of atoms that have internal structures that dictate their chemical and physical behavior.

Targets:• Describe the arrangement of elements in the periodic table in

order of increasing atomic number.• Distinguish between the terms group and period.• Apply the relationship between the electron arrangement of

elements and their position in the periodic table.• Apply the relationship between the number of electrons in the

highest occupied energy level for an element and its position in the periodic table.

• Discuss the similarities and differences in the chemical properties of elements in the same group.

• Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities.

Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Development of the Periodic Table

• Johan Dobereiner Grouped similar elements into groups of 3 (triads) such as chlorine, bromine, and iodine. (1817-1829).

• John Newlands

Found every eighth element (arranged by atomic weight) showed similar properties. Law of Octaves (1863).

• Dmitri Mendeleev Arranged elements by similar properties but left blanks for undiscovered elements (1869).

Describe the arrangement of elements in the periodic table in order of increasing atomic number.Distinguish between the terms group and period.

Development of the Periodic Table

• Henry Mosley Arranged the elements by increasing atomic number instead of mass (1913)

• Glen Seaborg Discovered the transuranium elements (93-102) and added the actinide and lanthanide series (1945)

Describe the arrangement of elements in the periodic table in order of increasing atomic number.Distinguish between the terms group and period.

Development of the Periodic Table

Elements are arranged by increasing

atomic number into periods (rows)

and groups or families (columns)

Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Arrangement of the Periodic Table• Metals

– Left side of the periodic table (except hydrogen).

– High electrical conductivity, high luster, ductile, malleable

– Alkali metals: Group 1 (1A)– Alkaline earth metals: Group

2 (2A)– Transition metals: Group B,

lanthanide & actinide series

Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Arrangement of the Periodic Table

• Nonmetals– Right side of the periodic

table– Poor conductors and

nonlustrous– Halogens: Group 17 (7A)– Noble gases: Group 18

(0)•

Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Arrangement of the Periodic Table

• Metalloids– Between metals and

nonmetals– Properties

intermediate between metals and nonmetals

Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table: pg 392-393

• Noble Gases: Outermost s and p sublevels are filled. – Ending configuration is

s2p6 (except He)– Eight valence electrons

(except He)– Row number equals

highest energy level

Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table: pg 392-393

• Representative Elements: Outermost s and p sublevels are partially filled.– Group 1,2, 13-17– 1 (s1); 2 (s2); 13 (s2p1); 14

(s2p2)…– Group number equals

valence electrons (1-2, 13-17 subtract 10)

– Row number equals highest energy level

Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table: pg 392-393

• Transition Metals– Filling the d & f sublevels

Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Shortcut Electron Configuration

Based on the electron configuration of the noble gases. He ends in 1s2; Ne ends in 2p6; Ar ends in 3p6; Kr ends in 4p6; etc.

• Write the electron configuration and orbital filling diagram for Se– Se has 34 electrons– Go back to the previous noble gas: Ar (18 electrons). Begin the

configuration with [Ar] which accounts for 18 electrons and then begin with 4s2. Continue until you reach 34 electrons

– [Ar]4s23d104p4

– [Ar] __ __ __ __ __ __ __ __ __ 4s 3d 4p

Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Shortcut Electron Configuration

• Write the electron configuration and orbital filling diagram for Au– Au has 79 electrons– Go back to the previous noble gas: Xe (54 electrons).

Begin the configuration with [Xe] which accounts for 54 electrons and then begin with 6s2. Continue until you reach 79 electrons

– [Xe]6s24f145d9

– [Xe] __ __ __ __ __ __ __ __ __ __ __ __ __ 6s 4f 5d

Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.

Shortcut Electron Configuration

Electron dot diagrams

Group 1A: 1 dot X Group 5A: 5 dots X

Group 2A: 2 dots X Group 6A: 6 dots X

Group 3A: 3 dots X Group 7A: 7 dots X

Group 4A: 4 dots X Group 0: 8 dots (except He) X

Discuss the similarities and differences in the chemical properties of elements in

the same group.

Group 1: Alkali Metals• Have 1 valence electron• Shiny, silvery, soft metals• React with water & halogens• Oxidize easily (lose electrons)• Reactivity increases down the

group

Discuss the similarities and differences in the chemical properties of elements in

the same group.

Group 17: Halogens • Have 7 valence electrons• Colored gas (F2, Cl2); liquid

(Br2);

Solid (I2)

• Oxidizer (gain electrons)• Reactivity decreases down the

group

Describe and explain the group and periodic trends in atomic radii, first

ionization energies and electronegativities.

Atomic RadiiSize of the atom

• Group trend– Atomic size increases as you move down a group of the

periodic table.– Adding higher energy levels.

Describe and explain the group and periodic trends in atomic radii, first

ionization energies and electronegativities.

Atomic RadiiSize of the atom

• Periodic Trend– Atomic size decreases as you move from left to right

across a period.– Effect of increasing nuclear charge pulls the electrons

closer to the nucleus.

.

Describe and explain the group and periodic trends in atomic radii, first

ionization energies and electronegativities.

First Ionization EnergiesThe energy required to remove the first electron from a gaseous

atom.Second ionization removes the second electron and so on.

Can be used to predict ionic charges.

• Group trend– Generally decreases as you move down a group in

the periodic table – Since size increases down a group, the outermost

electron is farther away from the nucleus and is easier to remove.

Describe and explain the group and periodic trends in atomic radii, first

ionization energies and electronegativities.

First Ionization EnergiesThe energy required to remove the first electron from a gaseous

atom.Second ionization removes the second electron and so on.

Can be used to predict ionic charges.

• Periodic Trend– Increases as you move from left to right across a

period.– Effect of increasing nuclear charge makes it

harder to remove an electron.

Describe and explain the group and periodic trends in atomic radii, first

ionization energies and electronegativities.

ElectronegativityTendency for the atoms of the element to attract electrons

when theyare chemically combined with atoms of another element.

Helps predict the type of bonding (ionic/covalent).• Group trend

– Generally decreases as you move down a group in the periodic table.

– For metals, the lower the number the more reactive.

– For nonmetals, the higher the number the more reactive.

Describe and explain the group and periodic trends in atomic radii, first

ionization energies and electronegativities.

ElectronegativityTendency for the atoms of the element to attract electrons

when theyare chemically combined with atoms of another element.

Helps predict the type of bonding (ionic/covalent).

• Periodic Trend– Increases as you move from left to right across a

period.– Nonmetals have a greater attraction for electrons

than metals.