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Why is the periodic table shaped like it is and how are the elements arranged?. The Periodic Table - History. Two scientists, D i mitri Mendeleev (Russia) and Lothar Meyer (Germany) properties of elements did not change smoothly with increasing atomic mas.

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Why is the periodic table shaped like it is and how are the elements arranged?


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The Periodic Table - History

Two scientists, Dimitri Mendeleev (Russia) and Lothar Meyer (Germany)

properties of elements did not change smoothly with increasing atomic mas.

Instead the properties of the elements repeated periodically.

Periodic Law: the properties of the elements repeat periodically as the elements are arranged in order of increasing atomic number (# of protons)

This periodic law is used to form the Periodic Table


Ii the periodic table l.jpg

II. The Periodic Table

  • John Alexander Newlands

    • Arranged elements in order of increasing atomic masses

    • Noticed some properties recurring over and over again – he called this the periodic law

  • Dmitri Mendeleev

    • Published the periodic table of elements

    • Very confident as he left spaces empty – assumed those elements were not yet discovered


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Vertically into Groups

Horizontally Into Periods

Elements are arranged:


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Groups of Elements

  • Vertical columns on the periodic table

  • Similar physical properties

  • Similar chemical properties


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Why?

Why?


If you looked at one atom of every element in a group you would see l.jpg

If you looked at one atom of every element in a group you would see…

Each atom has the same number of electrons in it’s outermost shell.


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105

Db

107

Bh

Valence Electrons

for group A elements

sblock

pblock

0

+1

Charge on stable monatomic ion

± 4

+3

-3

-2

-1

ns2np6

ns1

+2

ns2np3

ns2np1

ns2np2

ns2np4

ns2np5

ns2

8A

1A

3A 4A 5A 6A 7A

2A

Group B

Period number gives shell (n)


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Elements are arranged according to atomic #

and e- configuration.

Li: 3 e-’s 1s22s1

Na: 11 e-’s 1s2 2s2 2p63s1

K: 19 e-’s 1s2 2s2 2p6 3s2 3p64s1

Paramagnetic or diamagnetic?


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Valence orbitals: outer shell orbitals

beyond the closest noble-gas configuration

Valence electrons: “the ones that can react” (located in the valence orbitals).

The other e-’s are called core electrons and don’t react.

2s2

3s2

4s2

5s2

6s2

7s2

Elements in a vertical row have the same number of valence electrons.


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  • The number of outer or “valence” electrons in an atom effects the way an atom bonds.

  • The way an atom bonds determines many properties of the element.

  • This is why elements within a group usually have similar properties.


If you looked at an atom from each element in a period l.jpg

If you looked at an atom from each element in a period

you would see…


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The elements in the same horizontal row are called a period

Details of the Periodic Table - Period

Hydrogen and Helium are in period 1

Lithium through Neon are in period 2

4


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Periods on the Periodic Table

1

2

3

4

5

6


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Each atom has the same number of electron holding shells.

An example…


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The period 4 atoms each have 4 electroncontaining shells

4th Shell

K (Potassium)

Atom

Kr (Krypton)

Atom

Fe (Iron) Atom


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Sublevel Blocks

s1 s2p1 p2 p3 p4 p5 p6

1

2

3d1 - d10

4

5

6

f1 - f14


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A periodic table illustrating the building-up order.


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Electronic Structure

sblock

pblock

dblock

105

Db

107

Bh

fblock


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Condensed Ground-State Electron Configurations in the First Three Periods


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Los Alamos National Laboratory's Chemistry Division

Periodic Table of the Elements

noble

gases

non-metals

+1

metals

+2

+3

-3

-2

-1

typical charge on ion in binary compound

atomic number

atomic mass

transition metals

http://pearl1.lanl.gov/periodic/default.htm


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Discovery dates of elements


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Physical States at Room Conditions

gases

liquids

solids

105

Db

107

Bh

Modified from http://www.cem.msu.edu/~djm/cem384/ptable.html


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Natural vs. Man-Made Elements

natural elements

man-made elements

not found on Earth

105

Db

107

Bh


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MolecularNature of Free Elements

diatomic

tetratomic

octatomic

105

Db

107

Bh

All other free elements are atomic in nature.


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All Isotopes Radioactive

no stable forms

105

Db

107

Bh


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105

Db

107

Bh

Hydride Chemistry

(binary hydrogen compounds)

ionic

covalent

polymeric

salts with H-

molecules with +1

chains with H-to-H bonds

8A

1A

3A 4A 5A 6A 7A

2A

metallic

interstitial H2

Group B

mostly gases, some liquids

solids


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Important Groups


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Each group has distinct properties

  • The periodic Table is divided into several groups based on the properties of different atoms.


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Groups of Elements

  • Alkali Metals:

    • Group 1 metals

    • Soft, silver coloured metals that react violently with H2O to form basic solutions

    • Most reactive: cesium & francium


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Alkali Metals reacting with water:

  • Li (Lithium)

  • Na (Sodium)

  • K (Potassium)

  • Rb (Rubidium)

  • Cs (Cesium)

What would you expect from Francium?!?!


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Chemical Periodicity


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Potassium (K), in Water (H2O)

Alkali metals

Soft, silvery colored metals

Very reactive!!!


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Alkali Metal Family

Li

K

Na


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Alkaline Earth Metals

Silvery-White Metals

Fairly reactive

Many are found in rocks in the earth’s crust


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Halogens:

  • Group VII A, non-metals, highly reactive.

  • Fluorine is the most reactive

Noble Gases:

-Group VIII A

- Generally unreactive


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TheHalogenFamily

Cl2(g)

I2(s)

Br2(l)


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Chlorine (Cl) Bromine (Br) and Iodine (I)

Halogens

Most are Poisonous

Fairly reactive

Chlorine Gas was used as a chemical weapon during World War I.

It was used by the Nazis in World War II.


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Jellyfish lamps made with noble gases


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Colors Noble Gases produce in lamp tubes(discharge tubes):

  • Ne (Neon): orange-red

  • Hg (Mercury): light blue

  • Ar (Argon): pale lavender

  • He (Helium): pale peach

  • Kr (Krypton):pale silver

  • Xe (Xenon): pale, deep blue


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  • Transition Metals:

    Most are good

    Conductors of heat and electricity

    Ductile and malleable

    (easily bent/hammered into wires or sheets)

.Metalloids:

Metals very close to the “staircase” line

They have properties of metals and non-metals.

  • Si (Silicon) and Ge

  • (Germanium) are very

  • important “semi-conductors”


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How many things can you think of that have Transition Metals in them?


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What are semiconductors used in?


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Nonmetals

.Brittle

.Do not conduct heat and electricity


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Modern Periodic Table


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  • Columns are called Groups or Families

    •  Elements with similar chemical and physical properties are in the same column

  • Rows are called Periods

    •  Each period shows the pattern of properties repeated in the next period


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Main Group (Representative Group) - Groups IA - VIIIA

Transition Metals - Groups IB – VIIIB

Rare Earth Elements -

Lanthanides (Ce - Lu) and

Actinides (Th - Lr)


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Metals

about 75% of all the elements

lustrous, malleable, ductile, conduct heat and electricity

Nonmetals

dull, brittle, insulators

Metalloids

also know as semi-metals

some properties of both metals & nonmetals


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Atomic Size


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Smaller

Smaller

Li Be B C N O F

Na

WHY?

K


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Ionization Energy (IE) - The amount of energy needed to remove an electron from an atom or ion. Each electron in any atom or ion has a specific ionization energy.

First Ionization Energy - The amount of energy needed to remove an electron from the outermost shell of a neutral (uncharged) atom.


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The ionization energy is the energy that must be supplied to an atom in the gas phase in order to remove an electron. If the electron is the first one to be removed, one then refers to the first ionization energy for that element.

The first ionisation energy of sodium is 494 kJ.mol-1.

The ionisation energy tells us how easy it is to convert an element to a cation. The lower the first ionisation energy, the easier it is to convert an element such as sodium, Na, to its cation Na+.


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Ionization Energy:

The energy required to completely remove an e- from an atom in its gaseous state.

Mg(g) Mg1+ + e-

1st ionization energy

Mg1+(g) Mg2+ + e-

2nd ionization energy

Question: Which of the above ionizations would have the highest ionization energy and why?


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electron being lost:

1st 2nd 3rd 4th 5th 6th 7th


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The first ionisation energy of most elements are known to a good degree of accuracy. It is interesting to see how these values vary with atomic number Z:

The rare gases (He, Ne, Ar, Kr, Xe, Rn) appear at peak values of ionization energy, which reflect their chemical inertness, while the alkali metals (Li, Na, K, Rb, Cs) appear at minimum values of ionization energy, in keeping with their reactivity and ease of cation formation.


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Increases

Increases


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Ionization energy decreases as you go down a group

Ionization energy increases as you go from left to right in a period


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What is meant by metallic character?


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Periodicity of ionic radii:

The size of ions differ markedly from the size of their parent atoms. Take the case of sodium as an example:

The sodium atom (which has a single 3s electron), has an atomic radius of 186 pm. Upon taking up energy (the ionization energy), it loses this electron and is converted to the cation Na+, whose radius is 97 pm.


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The radius of cations is always smaller than the radius of the atoms from which they are derived, as shown in the figure on the right, which applies to the Group I elements:

The main reason for this is that whenever metals are converted to their cations, they always do so by losing the electrons in their highest energy level.

Further, since the ion has less electrons than the atom from which they are derived, there is less mutual repulsion between these electrons, and the electron orbitals shrink to some extent.


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What happens in the case of anions? Let's take the case of chlorine:

The chlorine atom has a covalent radius of 99 pm. It can gain an extra 3p electron (with release of energy, the electron affinity) in order to form the chloride anion Cl-, with ionic radius 181 pm.


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Periodicity of electronegativities:

When a covalent bond is formed between two identical atoms, such as H-H or Cl-Cl, the pair of electrons which joins the atoms is evenly shared between the two atoms. However, when two different atoms are joined together by a covalent bond, the sharing of the electron pair is not even, and the pair of electrons is shifted towards one of the atoms:

The covalent bond is said to have been polarized, and one refers to such a bond as a polar covalent bond.


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The ability to polarize a covalent bond differs from one atom to another, and is known as the electronegativity of the atom.

Electronegativities are measured on an arbitrary scale ranging from 0 (He) to 4.1 (F).

The rare gases (He, Ne, Ar, Kr, Xe, Rn) have zero electronegativities, i.e., they only form covalent bonds, and this only in exceptional cases, and have no tendency to attract electrons of that bond.

The halogens ( F, Cl, Br, I, At) appear at peak values of electronegativities. Within a period, they have the highest tendency to polarize covalent bonds. Note that within the group, the electronegativity decreases.

On the whole, electronegativities increase from left to right along a period (see Li to F) and decrease from top to bottom within a group (see F to At).


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Periodicity of electron affinity:

The electron affinity is the energy that is released when an atom in the gas phase gains an electron and is thus converted to an anion, also in the gas phase:

The electron affinity of chlorine is 349 kJ.mol-1.


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Electron affinities are difficult to measure and there is no reliable data available for most elements. However, the larger the atom, the lower its electron affinity, as shown with Group VII elements:

For reasons outside the scope of this discussion, the electron affinity of fluorine is an exception to this trend.


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Common Oxidation states: note the vertical similarities.


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