Periodic Table and Periodic Trends

Activity

•You and partner pull out all the writing utensils you have

•Create a method of organizing your writing utensils without talking to other groups

I. Development of Modern Periodic Table A. History

Lavoisier

• Antoine Lavoisier was first to compile list of known elements in 1790’s

Mendeleev

• Russian chemist, Mendeleev, organized a table by arranging elements in order of increasing atomic mass

• Mendeleev is credited with the first periodic table

• Mendeleev’s table predicted the existence and properties of undiscovered elements

Moseley• English chemist,

Moseley, discovered that each element contained a unique number of protons

• Moseley arranged elements in a table in order of increasing atomic number

Periodic Law:

There is a periodic repetition of chemical and physical

properties of the elements when arranged by increasing

atomic number

B. Modern Periodic Table

118

?

•Periodic Table of the Elements: organization of elements. Each square shows the name of an element, its chemical symbol, atomic number, and average atomic mass

Periodic Table

•elements arranged by increasing atomic number

Groups

•vertical columns•also called ______________

or just families (we’ll call them groups)

chemical families

Main Group•Groups 1,2 through 13-18 are

called main group elements (A groups )

Transition Metals

•Groups 3-12 are called transition elements (B groups )

•contains elements with similar chemical propertiesEX: Li, Na, K

Periods

•horizontal rows•physical and chemical

properties change somewhat regularly across a period

•elements close to each other in the same period are more similar than those further apart

•EX: K, Ca, Sc

The Staircase

•The two sides of the periodic table can be divided into metals and non-metals by the ____________ linestaircase

•Non-metals are found to the right of the staircase

•all elements to the left of the staircase are considered metals (except hydrogen)

•elements that border the staircase are called metalloids Ex. Si, Ge

A look into Metals

•What state of matter are most of them in at room temperature?

•Do any metals look familiar to you?

Solids (except for Mercury – it’s a liquid)

Silver, Gold, Platinum, Lead, Tin

Metals

•at the left of the staircase and bottom two rows

•conduct heat and electricity easily

•most are solid at room temperature (Hg is a liquid at RT)

Metals

•exhibit malleability (can be hammered or rolled into thin sheets)

•high tensile strength (ability to resist breaking when pulled)

Physical Properties of Metals

• range from soft (sodium) to extremely hard (platinum)

• Physical properties of metals include: malleable, ductile, lustrous, and conductivity of heat and electricity

• Malleable: can be hammered into thin sheets

• Ductile: can be pulled into wires• Lustrous: shiny appearance

Metals – Metalloids – Non-Metals

A look into Non-Metals

•What state of matter are most of them in at room temperature?

•Do any non-metals look familiar to you?

Gases

But some are liquids and solids too

Oxygen, Hydrogen, Nitrogen

Non-Metals

•Can be solids, liquids, or gases

• toward right of periodic table•most are gases at RT

Non-Metals

•examples of gases: nitrogen, oxygen, hydrogen

•examples of solids: carbon, phosphorus

Non-Metals

•solids are typically brittle•poor conductors of heat and

electricity

Pure carbon

Metalloids

• along the stair step line that separates metals and non-metals

• have some characteristics of metals and some of non-metals

• all are solid at room temperature

Metalloids

• less malleable than metals•not as brittle as nonmetals•semiconductors of electricity •used in electronics

II. Electrons and the Periodic Table

A. Valence electrons

•Valence electrons are the electrons in the outer-most energy level in an atom

•atoms in same group have similar chemical properties because they have the same number of valence electrons

Energy Level Diagrams

Nucleus

2 e- max

8 e- max

8 e- max

8 e- max• # e- = # p+ in an atom

• Fills center levels first

•Electrons are dots

Turn to set II of your Study Packet

Valence Electrons

Valence Electrons

• the period (row) can indicate the energy level of the element’s valence electrons

• the Roman numeral for the main group (A group) elements indicates the number of valence electrons for that group (exception: Helium – 2 valence electrons only)

Valence electrons

B. Electron blocks• s-block consists of groups 1(IA)

and 2 (IIA) and the s orbitals are being filled

• p-block consists of groups 13 (IIIA) through 18 (VIIIA) and the p orbitals are being filled

• the d-block consists of groups 3 (IIIB) through 12 (IIB) and the d orbitals are being filled

Electron Blocks

• the f-block includes the Lanthanide series and the Actinide series and the f orbitals are being filled

S Pd

f

1

2

3

4

5

6

7

4

5

Energy levels

3

4

5

6

2

3

4

5

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S

III. Properties of Elements

A. Hydrogen

•Group 1/IA and has 1 valence electrons

•Only non-metal in group 1•has 3 naturally-occurring

isotopes

A. Hydrogen

•can act like a nonmetal and lose an e- or act like a metal and gain an e-

•very reactive

B. Alkali Metals link

•Group 1 or IA• 1 valence electrons• forms ions with a +1 charge

(cations)•Most reactive metals •only found combined with

other elements in nature

C. Alkaline Earth Metals

•Group IIA or 2 and all have 2 valence electrons

• forms ions with a 2+ charge (cations)

• less reactive than alkali metals, but still pretty reactive

D. Halogens

•Group VIIA or 17 and have 7 valence e-,

•most reactive nonmetals•only found combined with

other elements in nature• forms ions with 1- charge

(anions)

Halogens are commonly referred to as “Halides”

(write that down)

E. Noble Gases (Inert Gases)

•Group VIIIA or 8 •8 valence electrons, except

He that has 2 valence e-

• rarely reactive•most are gases at RT (room

temperature)

Noble Gases

Examples: helium, neon

F. Transition Metals

•Groups 3 through 12 •fills the d block

G. Inner Transition Metals

•Also referred to as rare earth metals

•The two rows on the bottom of the periodic table

• Include Lanthanide and Actinide Series

•Fills the f block

IV. Periodic TrendsKeep these 3 factors in mind when considering periodic trends:

1. Nuclear charge

•Whenever a proton is added to the nucleus, it creates a stronger “nuclear magnet” pulling in the electrons even more.

• Electrons added to the same energy level (period) will be pulled in tighter toward the nucleus_.

• Ex:

2. Nuclear Shielding

•When an energy level is added to the atom (each new period you are adding “layers” between the nucleus and the valence electrons.

• As energy levels are added, the atom becomes larger and you dilute the pull of the nucleus for the valence electrons because not only are there more “layers” , but the valence electrons are also now farther from the nucleus.

• (It is easier to remove a valence electron as energy levels are added.)

•Ex.

3. Octet Rule

• Atoms will lose, gain or share electrons so they can achieve the electron configuration of the closest noble gas.

• As elements get closer to the noble gases on the periodic table (further right ), the greater the attraction for electrons.

Noble gases DO NOT attract electrons

• Elements on the left side of the periodic table want to lose electrons, so they will not have a great attraction for electrons.

• EX

A. Atomic Radius

Atomic radius is defined as one-half the distance between nuclei of two like atoms in a diatomic molecule

Ex.

•measured in picometers (pm), 10-12 m or Angstroms (A), 10-10 m

•atomic radius indicates relative size of the atom

Just the main groups

1. Group trends.

•Atomic radius generally increase as you move down a group. This is mainly due to succeeding energy levels being filled.

2. Period trends

•Atomic radius generally decrease as you move across a period from left to right. This is mainly due to increasing nuclear charge

Atomic Radius

•Smallest atomic radius = Helium

•Largest atomic radius = Francium

Decreases

increases Increases

B. Ionic Radius

B Ionic Radius

• Ionic radius is the measurement of an ion in a crystal lattice

•The units of measurements is picometers (pm) or Angstrom (Å)

1. Group trends for ionic radiiIonic radius increase as you move

down a group. This is because of the added “layers” of electrons.

2. Period Trends for ionic radiiIonic radius decrease as you move

left to right across a period. This is mainly due to the nuclear charge.

Ionic Radius

Decreases

increases Increases

C. Ionization Energy

• Ionization energy is defined as the amount of energy required to remove an electron from an atom.

0

500

1000

1500

2000

2500

0 5 10 15 20Atomic Number

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1. Group Trends

• Ionization energy decrease as you move down a group. This is mainly due to the energy levels.

2. Period trends

• Ionization energy increases as you move left to right across a period. This is mainly due to octet

Ionization Energy

•Smallest Ionization NRG =Francium

•Largest Ionization NRG = Helium

Increases

decreases

Increases

D. Electronegativity

•Electronegativity is defined as the tendency for atoms of the element to attract electrons when they are chemically combined with atoms of another element.

1. Group trends

•Electronegativity decreases as you move down a group. This is because the distance from the nucleus is increasing

2. Period Trends

• Electronegativity increases as you move left to right across a period. This is mainly due to the octet rule.

Note: Noble gases have no electronegativity because they don’t attract electrons at all. Again, think of the octet rule.

Electronegativity

•Most Electronegative = F•Least Electronegative = Fr

(+ all noble gases, of course)

Increases

Decreases

Increases

decreases

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increasing

increasing

decr

ease

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incr

ease

sincreasing

incr

ease

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decreases

incr

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Atomic Radius

Ionic Radius

Ionizing Energy

Electronegativity

Nuclear Charge

Shielding

decreases

Ato

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Rad

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Ion

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Ion

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Ele

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Nu

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