Introduction to the Periodic Table

Introduction to the Periodic Table

alkaliAlkaline earth

<---transition-------------------->

<-----------------Inner transition------------------>

halogens

Noble gas

I. History of the Periodic TableDemitri Mendeleev (1860’s

Russia)• Arranged known

elements:– by mass– in vertical columns by

Phs/Chem Prop.• Left blank spaces

(very accurate!)• Published his PT in 1871

Henry Mosley (1913)• Found a way to determine the atomic

number of elements using their emission frequency

• Arranged the PT by atomic number• “Father” of the modern periodic table

As time passed:*more elements were discovered*when placed by P/C properties, masses

were out of order

II. The Modern Periodic Table• Arises from Periodic

Law:A. Includes:

1. periods/rows2. groups/families

• There are 4 categories of groups:

1. Representative (also Main Group Elements)

2. Transition elements 3. Inner transition (also

Lanthanide and Actinide) 4. Noble Gases

• Some groups have special names: – 1A (Alkali Metals)– 2A (Alkaline Earth Metals)– 7A (Halogens)– 8A I (Noble Gases)

• The letters A and B in the group distinguish families – A = representative – B = transition

Properties and TrendsElectron Configurations

Metallic Character

i) L to R: metals to nonmetals

ii) Nonmetals are at the right of the Table. They tend to be insulators and react easily with metals.

iii) Metalloids separate the metals and nonmetals and have intermediate properties

iv) Noble Gases exist at the extreme right, are chemically stable and have full valence shells

Atomic Radius– Atomic radius – half distance between the nuclei of two

atoms of same element• Shielding effect – inner energy levels ‘shield’ the outermost

electrons from the positive charge pull of the nucleus– Atomic radius increases as you move down the groups• Great distance (adding energy levels) from nucleus = less pull

towards center– Decreases as you move left to right• More pull from nucleus (more protons), but no new distance

– EXCEPTION: Noble Gases – much bigger than group 17 – full outer shell

Atomic Radii

Ionization Energy– Ion – atom which has gained or lost electrons• Cation – (+) charged ion (lost e-)• Anion – (-) charged ion (gained e-)

– Ionization energy – the energy that is required to remove an e- from an atom

– Decreases as you move down the periodic table• Outermost electron gets further from nucleus, easier to pull off

– Increases as you move left to right• No more distance from nucleus, but higher charge = held more

tightly

Ionization Energy

Electron Affinity– Electron Affinity – the energy change that occurs when

an electron is ADDED to a neutral atom• Bigger negative number = easier to add e-

– Harder to add as you go down a group• Further distance from nucleus & more inner e- = more

repulsion felt from e-

– Easier to add from left to right• Increased nuclear charge = more attraction to nucleus (+)• Noble Gases – don’t accept e-

– Halogens gain most easily, they want to complete that ‘perfect eight’

Electron Affinity

Electron Affinity & Ionization Energy

Ionization Energy vs. Electron Affinity

Ionic Size (Radius)

– Ionic Radius- ½ the distance between the nuclei of two ions

– Cations (+) are always SMALLER than neutral atom• Nuclear charge the same, less e- = strong pull inwards

– Anions (-) are always LARGER than neutral atom• Nuclear charge the same, MORE e- = less pull inwards

– Increases as you go down a group– Decreases from left to right

Ionic Radii

Electronegativity

–Electronegativity – tendency for an element to have a stronger pull on the shared e- in a covalent bond (values btwn 0-4)–Decreases down a group – less likely to

keep the shared e-

–Increases from left to right – more likely to have the shared e-

Electronegativity