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 elements2. Named groups3. 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 IIMoseley: 1914

Atomic numbereach 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 LawThe physical and chemical properties of the elements are periodic functions of their atomic numbers.

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

How to use the periodic table…

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

SymbolBasically 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 averageAlso notice the different locations of the atomic mass and atomic number.

Groups (families)The Columns Elements in groups have similar chemical properties

Periods The RowsElements properties vary across periodsThe 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 metalsd-block elements

• Lanthanides and actinidesf-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.

• Procedures1. 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 element3. 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. 1351: Identify period, block, group, element[Kr]5s22. 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 reactiveLess reactive 4<3<2<1 More reactiveLess reactive 4<5<6<7 More reactive

• Sample Problem D Solutiona. 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 TrendsEssential 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 periodsBecause of increasing positive charge of the nucleus Holds electrons more tightly

INCREASES down groupsHigher principle quantum numberValence 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) Ionization Energy (IE) DefinitionDefinition

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 Ionization Energy (IE)(IE)

IE increases across periods because of increasing nuclear charge. Why?

Higher charges 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 AffinityThe 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 AffinityPeriod Trends: Halogens (VII A) gain

electrons the most readilyGroup 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