Assignment 10/24

• Read Chapter 6, section 1. You have 10 minutes.

• Answer the questions below:– 1. How is the modern periodic table arranged?

Who organized elements with this arrangement?– 2. Who arranged the periodic table by increasing

atomic mass?– 3. What is the name of Group 18 elements?

11/8 Activity• 1. Write the electron configuration of your element.• 2. Name the group (if applicable) that your element falls in.

List all properties of that group. Be detailed!• 3. How many valence electrons does your element have in the

outer shell.• 4. Draw the Lewis dot structure for your element.• When finished…transfer all the information about the

elements and group to a piece of butcher paper. Show me your work and hang it on the wall somewhere. We will do a gallery walk of the groups at the end of the period today or beginning tomorrow.

11/11 Today you will need…

• A clean sheet of paper, something to write with, and the paper from the side table.

• We will do a gallery walk activity using the information from Friday about the groups on the periodic table.

• You will color code a periodic table when the gallery walk is complete

Gallery Walk…• When finished, answer the following questions on a clean

piece of paper and turn in:

– 1. Name Group 1, Group 2, Group 17, and Group 18 elements.– 2. Identify the properties of each of the groups listed in number 1.– 3. Determine the number of valence electrons in groups

1,2,13,14,15,16,17,18.– 4. Draw the electron configuration for ONE element from each group

listed in number 1.

Chapter 5-The Periodic Law

5.1-History of the Periodic Table5.2-Electron Configuration & the Periodic Table

5.3-Electron Configuration & Periodic Properties

5.1-History of the Periodic Table

Pages 123-127

Mendeleev

• Dmitri Mendeleev (1869, Russian)–Organized elements

by increasing atomic mass.

–Elements with similar properties were grouped together.

–There were some discrepancies.

Mendeleev

• Dmitri Mendeleev (1869, Russian)–Predicted properties of undiscovered elements.

Moseley

• Henry Moseley (1913, British)–Organized elements by increasing atomic number.

–Resolved discrepancies in Mendeleev’s arrangement.

–Periodic Law-the physical and chemical properties of the elements are periodic functions of the atomic numbers.

Organization of the Elements

• 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.

Additions to Mendeleev’s Periodic Table

• Noble gases– Group 18– Argon discovered in 1894– Took so long to discover because very unreactive

• Lanthanides– 14 elements with atomic numbers from 58-71– Placed below the periodic table to conserve space

• Actinides– 14 elements with atomic numbers 90-103– Also placed below periodic table

Warm-Up 11/11

• Define Periodic Law.

• Where are the Lanthanides and Actinides located on the periodic table.

• Describe, specifically what contributions Mendeleev and Moseley made towards the creation of the periodic table. (i.e dates, arrangement, etc.)

5.2-Electron Configuration & the Periodic Table

Pages 128-139

Periods & Blocks of the Periodic Table

• Length of period (row) determined by how many electrons can occupy the sublevels being filled.– 1st period-1s sublevel being filled with 2 electrons

2 elements, H & He– 3rd period-3s & 3 p sublevels being filled with 2+6

electrons 8 elements• Periodic table is divided into “blocks” based

on the filling of sublevels with electrons.

Blocks of the Periodic Table

Determining Period from Configuration

• An element’s period can be determined by looking at its electron configuration

• The highest occupied energy level corresponds to the element’s period– As: [Ar]3d104s24p3

• 4 in 4p3 indicates that the highest energy level that electrons occupy is the 4th. Therefore, As is located in the 4th period of the periodic table.

Metallic Character

• Metals• Nonmetals• Metalloids

Areas of the Periodic Table

• Main Group Elements• Transition Metals• Inner Transition Metals

s-Block Elements: Groups 1 & 2

• Chemically reactive metals• Include the alkali metals and the alkaline earth

metals

Alkali metals

• Group 1 metals• ns1

• Silvery appearance and very soft• Not found pure naturally because so reactive• Because of extreme reactivity with moisture,

usually stored under kerosene• Video: Disposal of Surplus Sodium• Video: Alkali Metals in Water

Alkaline-Earth metals

• Group 2 metals• ns2

• Harder, denser, & stronger than alkali metals• Also too reactive to be found free in nature

(but less reactive than Gp. 1)• Video:

Magnesium/silver nitrate mixture reacting with water

d-Block Elements: Groups 3-12

• Metals with typical metallic properties• Called “transition elements”• Typically less reactive than Gps. 1&2, & some are extremely

unreactive• d sublevels first appears at the 3rd energy level• Fills after 4s• Variations from expected in d-block, so elements in the same

group do not necessarily have the same outer e- configuration

p-Block Elements: Groups 13-18

• p and s-block elements together called “main-group elements”

• Total number of electrons in highest energy level=group # - 10– Group 17 elements have 17-10=7 outer “valence”

electrons• Properties of p-block elements vary greatly

since metals, nonmetals, and metalloids are contained here

p-block Elements

• Halogens– Group 17 nonmetals– Most reactive nonmetals

• React with most metals to form salts

• Metalloids– Fall on both sides of a “stair-step” line separating

metals and nonmetals– Semi-conductors

f-Block Elements: Lanthanides & Actinides

• Lanthanides– Top row of f-block– 14 elements– Shiny metals similar in reactivity to the alkaline-earth

metals• Actinides

– Bottom row of f-block– 14 elements– All radioactive– 1st 4 elements found naturally on Earth; remainder

only lab-made elements

5.3-Electron Configuration & Periodic Properties

Pages 140-154

Remember the Periodic Law

• When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

0

50

100

150

200

250

0 5 10 15 20

Ato

mic

Ra

diu

s (

pm

)

Atomic Number

Warm Up 11-12

Explain the joke:

Yo momma so fat if she was on Mendeleev’s periodic table she would be at the end.

1

2

3

4 5

6

7

½ the distance between the nuclei of identical atoms that are bonded together

Atomic Radius

0

50

100

150

200

250

0 5 10 15 20Atomic Number

Ato

mic

Ra

diu

s (

pm

)Atomic Radius

Li

ArNe

KNa

Ionization Energy• First Ionization Energy- the energy required to

remove one electron from a neutral atom

1

2

3

4 5

6

7

Ionization Energy

• First Ionization Energy

0

500

1000

1500

2000

2500

0 5 10 15 20Atomic Number

1s

t Io

niz

ati

on

En

erg

y (k

J)

KNaLi

Ar

NeHe

Ionization Energy• Why opposite of atomic radius?

–In small atoms, e- are close to the nucleus where the attraction is stronger

• Why small jumps within each group?

–Stable e- configurations don’t want to lose e-

Ionization Energy

• Successive Ionization Energies

Mg 1st I.E. 736 kJ

2nd I.E. 1,445 kJ

Core e- 3rd I.E. 7,730 kJ

Large jump in I.E. occurs when a CORE e- is removed.

Al 1st I.E. 577 kJ

2nd I.E. 1,815 kJ

3rd I.E. 2,740 kJ

Core e- 4th I.E. 11,600 kJ

Ionization Energy

• Successive Ionization Energies Large jump in I.E. occurs when a CORE e- is

removed.

Electron Affinity

• Energy change that occurs when an electron is acquired by a neutral atom

• Tends to become less negative (less energy released) DOWN and to the LEFT

1

2

3

4 5

6

7

Electronegativity

• A measure of the ability of an atom in a chemical compound to attract electrons

• Most electronegative element is fluorine– Given arbitrary value of 4; all others relative

1

2

3

4 5

6

7

AC

B

D

EF

G

Label the Groups:

Atomic RadiusQ: Why larger going down?

A: Higher energy levels have larger orbitals

Q: Why smaller to the right?

A: Increased nuclear charge without additional energy levels to shield the electrons pulls them in tighter.

Q: Why is Ionization Energy greatest in upper right corner?

A: In small atoms, e- are close to the nucleus where the attraction is stronger. More protons = more positive charge.

Ionization Energy

Ionic Radius: The distance from the nucleus to the outer orbital of an ion.

Ions Atoms with an electric charge

Cations Anions

+ charge- charge

Lose electrons Gain

Electrons

SmallerBigger

Ionic Radius

Q: Why are cations smaller than the neutral atom?

A: When cations lose their electrons they also lose an energy level.

Q: Why are anions larger than the neutral atom?

A: As they gain electrons there is more repulsion between them, which pushes them farther from the nucleus.

Examples

• Which atom has the larger radius?

Be or Ba

Ca or Br

Examples

• Which atom has the larger radius?

Be or Ba

Ca or Br

Ba

Ca

Examples

• Which atom has the higher 1st I.E.?

N or Bi

Ba or Ne

Examples

• Which atom has the higher 1st I.E.?

N or Bi

Ba or Ne

N

Ne

Examples

• Which has the greater electonegativity?

K or Li

Al or Cl

Examples

• Which has the greater electonegativity?

K or Li

Al or Cl

Li

Cl

Examples

• Which particle has the larger radius?

S or S2-

Al or Al3+

Examples

• Which particle has the larger radius?

S or S2-

Al or Al3+

S2-

Al