# F321 - Topic 1 Atoms and Electron Structure

• Created by: Sophie
• Created on: 26-01-14 14:18

## Atomic Structure

Atomic Structure

• The nucleus contains almost all of the mass of an atom, as it cotains protons and neutrons.
• The nucleus of an atom contains all of the positive charge.
• The electrons orbit outside the nucleus.
• Electrons are negatively charged.

Definitions

• Atomic number = number of protons in the nucleus.
• Mass number = number of protons and neutrons in the nucleus.
• Number of neutrons = mass number - atomic number.
• Number of electrons = number of protons (in a neutral atom)
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## Isotopes and Ions

Isotopes

• Isotopes are atoms with the same number of protons but different number of neutrons.
• Isotopes of an element have the same chemical properties, as they have the same electron arangement.

Ions

• Ions are formed when atoms gain or lose elctrons.
• E.g. An atom Cl has 17 electrons, a Cl- ion has gained one electron.
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## Relative Atomic Mass and Relative Isotopic Mass

Relative Atomic Mass

• Average mass of an atom relative to 1/12 of the mass of a carbon-12 atom.
• To calculate relative atomic mass, add together (mass number x percentage/100).

Example

• 75% of Cl atoms have a mass number of 35
• 25% of Cl atoms have a mass number of 37
• Average mass of a Cl atom    = (mass no x percent/100) + (mass no x percent/100)

= (35 x 75/100) + (37 x 25/100)

= 35.5

Relative Isotopic Mass

• Mass of a particular isotope relative to 1/12 of the mass of a carbon-12 atom.
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## Orbitals

Orbitals

• An orbital is a region that can hold up to two electrons with opposite spins.
• Orbitals have different shapes called - s, p, d and f.

S orbitals

•  S orbitals are spherical in shape and come in sets of one.

P orbitals

• P orbitals are hour - glass shaped and come in stes of three.

D orbitals

• D orbitals come in sets of five which can hold up to 10 electrons.
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## Energy Levels

E.g.    Iron - 26 electrons - 1s2 2s2 2p6 3s2 3p6 4s2 3d6

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## Successive Ionisation Energies

Successive Ionisation Energies

• Energy change when one mole of electrons is removed from one mole of gaseous positively charged ions.
• In an element, successive ionisation energies get bigger because the emaining electrons are held more tightly by the unchanged nuclear charge.

First Ionisation Energy

• Energy change when one mole of electrons is removed from one mole of gaseous atoms.
• M(g) -----> M+(g) + e-
• Ionisation is easier if :
• The nuclear charge is smaller.
• The electron is further away from the nucleus.
• There is more shielding.
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## First Ionisation Energies - First 20 Elements

• First ionisation energy shows periodicity.

Factors affecting the size of ionisation energy

• The size of that attraction will be governed by:
• The charge on the nucleus.
• The distance of the electron from the nucleus.
• The number of electrons between the outer electrons and the nucleus (shielding).
• Whether the electron is on its own in an orbital ( paired electrons are easier to remove).
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## The Pattern in Periods 2 and 3

• Why the ionisation energy increases:
• There is an increasing number of protons in the nuceus.
• The increasing nuclear charge, drags the outer electrons in closer to the nucleus.
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## Trends in ionisation energy down a group

• Why:
• There is a bigger number of protons.
• But, the shielding and nuclear radius overcomes the nuclear charge.
• E.g.
• Li -1s22s11st I.E. = 519 kJ mol-1
• Na -1s22s22p63s11st I.E. = 494 kJ mol-1
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## Blocks in the Periodic Table

Blocks

• The s block is group 1 and 2.
• The p block is group 3 and 8.
• The d block is between the s and p blocks.
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