Ch. 5: Periodic

Table

C. Goodman, Doral Academy Preparatory High School, 2011-2013

Essential Question: Section 5.1

1. What is the history of the development of the Periodic Table?

2. What is the periodic law, and how can it be used to predict physical and chemical properties of elements?

3. What is the overall organization of the modern Periodic Table? 1. Three types of elements

2. Named groups

3. Other families

Section 5.1 Vocabulary

•Mendeleev

• Moseley

• Periodic Law

• Period

• Group

•Main group elements

Periodic Table - Definition

Periodic Table -- an arrangement of the elements in order of their atomic numbers so that elements with the same chemical properties are in the same group (family). Examples: halogens, noble gases, alkali metals.

Why is it cool?

• http://www.youtube.com/watch?v=u2ogMUDBaf4&playnext=1&list=PLAC3A0775D813045F&feature=results_main

History of the Periodic Table I

• Mendeleev: 1869

– Atomic mass

– Repeating (periodic) patterns of reactivity

•In his favor: predicted

the discovery of Gallium,

which was isolated in his

lifetime

•Certain characteristic

properties of elements

can be foretold from their

atomic weights

•Problem:

Iodine and Tellurium

History of the Periodic Table II

Moseley: 1914 Atomic number

each element has a unique atomic number; resequenced the table by electronic charge (=atomic #) rather than atomic weight.

Periodic Law

•In his favor:

solved the “Iodine and

Tellurium” problem

Moseley’s Periodic Law - Definition

• Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic numbers.

http://www.youtube.com/watch?v=OduTDU

GeAXEFind

How to use the periodic table…

Atomic number: # of protons in the nucleus of an atom

Symbol Basically the abbreviation for the element

Average atomic mass # of Protons + # of Neutrons (amu) Weighted average mass of isotopes of the element

Remember nuclear notation for isotopes? Notice that the atomic mass is a whole number – it’s not an average Also notice the different locations of the atomic mass and atomic number.

Groups (families)

The Columns

Elements in

groups have

similar chemical

properties

Periods

The Rows

Elements properties vary

across periods

The length of each period is

determined by the number of

electrons that can occupy the

sublevels being filled in that

period

Important terms

• Main Group Elements s-block + p-block elements

• Transition metals d-block elements

• Lanthanides and actinides f-block elements

• Metalloids, metals, non-metals (see below)

2 Main Sections in Periodic Table

Metals Nonmetals

- Majority of elements -Good Electrical & Heat Conductors - Room temperature = most solids -Contain properties

- Malleability - Ductility - High tensile strength

- Poor Electrical Conductors - Poor Heat Conductors - Room temperature = most gases - One is a liquid at r.t. = Bromine - Solid nonmetals generally brittle

Names of groups

• Group 1a – Alkali metals

• Group 2a – Alkali earth metals

• Group 7a (17) – halogens

• Group 8a (18) – Noble gases

Names of families– Transition metals

Names of families– Semiconductors

Section 2: Electron Configuration

1. What is the relationship between the location of atoms in a group, their electron configuration, and their chemical and physical properties?

2. What are the s-, p-, d-, and f-blocks, and how can their electron configurations of their elements be determined?

Section 5.2 vocabulary

• Ion

• Valence

• Valence electrons

• s-, p-, d- and f-block elements

Valence Electrons • Valence = outermost

energy level in which contains electrons (in unexcited state).

• Valence electrons are the electrons on the outermost energy level of the element.

• The number of valence electrons determines the type of chemical reactions available to the element!

What does this have to do with groups?

• All main group elements in a particular group have the same number of valence electrons.

• Prove it?

• Heh heh heh – that’s your job!

• Write electron configurations, noble gas notation, of the s and p block elements in the first 5 rows. Write valence electrons (s/b only) in contrasting color

• Elements hydrogen – xenon, for columns #1a, 2a, 3a, 4a, 5a, 6a, 7a, 8a

• Huhhh? See next slide

Valence electrons &energy levels

• Purpose: to determine the relationship between the group number and number of valence electrons.

• Procedures 1. For each of the elements in the first 5 periods of the following

groups… group (1a, 2a, 3a, 4a, 5a, 6a, 7a, and 8a)

2. Write the name of the element

3. Write the type of element

4. Write the electron configuration, using noble gas notation

• Conclusion (Answer the following question)s: – 1. How does the group # relate to the number of valence

electrons?

– 2. How do you think the chemical reactivity of the elements in a particular group, relates to this number of valence electrons?

Valence electrons & Energy Levels

• Conclusion (Answer the following questions): – 1. How does the group # relate to the number of valence

electrons?

– 2. How do you think the chemical reactivity of the elements in a particular group, relates to this number of valence electrons?

Group 14

Element Electron configuration – noble gas notation

Number of valence electrons

Carbon [He]2s22p2 4

Silicon [Ne]3s23p2 4

Germanium [Ar]4s23d104p2 4

Tin [Kr]5s24d105p2 4

Valence Electrons

• Main group elements have characteristic numbers of valence electrons.

• Group 1 – 1 valence electron

• Group 2 – 2 valence electrons

• Groups 13-18

– # valence electrons = Group # - 10

– Example: Group 13 elements have 13-10 = 3 valence electrons

Valence Electrons • Main group elements have characteristic

numbers of valence electrons.

• S block

– Group 1 – 1 valence electron

– Group 2 – 2 valence electrons

• P block

– Groups 13-18

• Group # - 10

• Example: Group 13 elements have 13-10 = 3 valence electrons

Relationship Between Periodicity and Electron Configurations

Sample problem A

1. Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe]6s2 is located.

2. Without looking at the periodic table, write the electron configuration for the Group 1 element in the third period. Is this element likely to be more reactive or less reactive than the element described in (a)?

Sample problem B

An element has the electron configuration [Kr] 5s24d5. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group.

Sample problem C

Without looking at the periodic table, write the outer electron configuration for the Group 14 element in the second period. Then, use your periodic table to name the element, and identify it as a metal, nonmetal, or metalloid.

In book, p. 135

1: Identify period, block, group, element

[Kr]5s2

2. write configuration of…

a. Group 2 elements

b. the group 2 element in the fourth period.

c. the element in the 3rd period, group 15

Sample Problem C Solution

• p-block (group # >12

• 14-10 = 4 electrons in s, p

• 2 e- in s, 2e- in p

• The outer electron configuration is 2s22p2.

• The element is carbon, C, which is a nonmetal.

More practice problems!

Name the block and group in which each of the following elements is located in the periodic table. Use the periodic table to name each element. Identify each element as a metal, nonmetal, or metalloid. Finally, describe whether each element has high reactivity or low reactivity.

1. [Xe] 6s24f145d8

2. [Ne]3s23p2

3. [Ne]3s23p5

4. [Xe]4f66s1

How do I know which groups are more

reactive than others?

For main group elements, look at the

number of valence (s/p) electrons

8 valence electrons (noble gases)

= not reactive

Less reactive 4<3<2<1 More reactive

Less reactive 4<5<6<7 More reactive

• Sample Problem D Solution a. The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially

filled with nine electrons. Therefore, this element is in the d block.

• The element is the transition metal platinum, Pt, which is in Group 10 and has a low reactivity.

• b. The incompletely filled p sublevel shows that this element is in the p block.

• A total of seven electrons are in the ns and np sublevels, so this element is in Group 17, the halogens.

• The element is chlorine, Cl, and is highly reactive.

Section 2 Electron Configuration

and the Periodic Table Chapter 5

Periods and Blocks of the Periodic Table, continued

• Sample Problem D Solution, continued • c. This element has a noble-gas configuration and thus is in Group 18

in the p block.

• The element is argon, Ar, which is an unreactive nonmetal and a noble gas.

• d. The incomplete 4f sublevel shows that the element is in the f block and is a lanthanide.

• Group numbers are not assigned to the f block.

• The element is samarium, Sm. All of the lanthanides are reactive metals.

Section 2 Electron Configuration

and the Periodic Table Chapter 5

Periods and Blocks of the Periodic Table, continued

Section 3: Periodic Trends Essential Questions

1. Compare the periodic trends of atomic radii,

ionization energy, electronegativity, and state the

reasons for these variations.

2. What are valence electrons, and how many are

present in atoms of each main-group element?

3. Compare the atomic radii, ionization energies, and

electronegativities of the d-block elements with

those of the main-group elements.

Section 5.3 Vocabulary

• Atomic radius

• Ion

• Ionization energy

• Cation

• Anion

• Electron affinity

• Electronegativity

Atomic Radii: ½ the distance between

the nuclei of identical atoms bonded

DECREASES across periods

Because of increasing positive

charge of the nucleus

Holds electrons more tightly

INCREASES down groups

Higher principle quantum

number

Valence electrons in higher

main energy levels

Located farther from the

nucleus

Periodic Trends: Atomic Radii

IONS

Watch it kid,

I’ve got my ion you.

Ion – Definition An Ion is…

An atom or group of bonded atoms, which has a positive (+) or negative (-) charge.

There are two kinds of ions…

Cation • CRUNCH (subtract) an electron

• This results in a positive charge

• When an electron is removed, the atom loses bulk (like a muscle which shrinks when it atrophies)

• So, the radius of a cation is smaller than the atomic radius

Anions

• ADD an electron

• This results in a NEGATIVE charge

• When an electron is added, the atom gains bulk (like a muscle which grows when you work out)

• So, the radius of a anion is larger than the atomic radius

Two sodium atoms bumped into each other.

One said: "Why do you look so sad?“ The other responded:

"I lost an electron.“

The first one asked "Are you sure?“

The other replied "I'm positive."

Which elements form

which type of ion?

Metal on left tend to form cations (+)

Nonmetals at the upper right tend to form

anions (-)

Hydrogen is a non-metal, but it forms a

cation (+)

Periodic Trends: Ionic Radii – very similar to trends for atomic radii

Ionization Energy (IE) Definition

Ionization energy is…

The energy required to remove one electron from a neutral atom of an element, forming a cation.

Note: does not apply to formation of anions!

Unit of measure: kJ/mol

It takes energy to “steal” electrons

Ionization energy is

(almost always)

positive (J)

Look p. 145, Figure 3.4

Low IE - lose electrons easily

High IE – it’s harder to lose electrons

Ionization Energy:

Atom (neutral) + energy

Cation (A+)+ (e-) (removed)

Ionization Energy (IE)

IE increases across periods because of increasing nuclear charge.

Why?

Higher nuclear positive charge attracts e- in same energy level more

strongly

Trend down groups decrease because of more electrons between the

nucleus and furthest out electrons (lower nuclear charge)

Watch out!!!

Ionization energy only applies to the

formation of cations.

If you want to talk anion formation,

you need electron affinity.

Electron Affinity The energy change that occurs when an

electron is acquired by a neutral atom

Most release energy when they acquire an

electron: Take a neutral atom (A) add an electron (e-)

get an anion (A-) + energy is released

Some must be “forced” to gain an electron

by adding energy: A + e- + energy A-

Common unit used is kJ/mol

Some positive affinities are difficult to

determine with any accuracy

Here is a mnemonic for electron affinity:

If the ion is negatively charged (anion), the electron affinity is more

strongly negative.

Why do some elements give up their electrons more easily?

Q: Which group of elements do you think hold on to their

electrons the most strongly?

A: The noble gases! Their valence is full, so it’s very hard to

remove their electrons!

Electron Affinity Period Trends: Halogens (VII A) gain

electrons the most readily

Group Trends: Electrons add with greater

difficulty down a group

Electron Affinity

• Question Based on this trend what elements will most likely form cations?

• Which will most likely form anions?

• Which will most likely have small ionic radii?

Adding Electrons to Negative Ions

Difficult to add a second electron

to an already negatively charged ion

Second electron affinities are

therefore all positive

Halogens become negative ions by

adding one electron (i.e. Cl-)

Electronegativity

• A measure of the ability of an atom in a

chemical compound to attract

electrons from another atom in the

compound

• Most electronegative element is

Fluorine

Electronegativity

Trend across periods increase (some exceptions)

More or less, trend down groups decrease

Similar to IE

Additional Help Website

http://www.hcc.mnscu.edu/chem/stacks.php