Chemistry - ATOMIC STRUCTURE AND THE PERODIC TABLE

Relative charges and masses of subatomic particles

Proton +1 mass 1, neutron 0 mass 1, electron -1 mass about 1/1836 (very small)

Where are subatomic particles found

Protons and neutrons in nucleus, electrons in shells (energy levels)

Atomic number

Number of protons (equals electrons in a neutral atom)

Mass number

Total number of protons and neutrons

How to calculate neutrons

Mass number minus atomic number

Isotopes

Atoms of the same element with different numbers of neutrons

Why isotopes have same chemical properties

Same number of electrons so same electron arrangement

Relative atomic mass (Ar)

Weighted mean mass of isotopes based on their abundance

Equation for relative atomic mass

Sum of (mass × abundance) divided by total abundance

Plum pudding model

A sphere of positive charge with electrons embedded in it

Alpha scattering experiment showed

Atom is mostly empty space with a small dense positive nucleus and electrons around it

Experiment that led to nuclear model

Alpha particle scattering experiment by Rutherford

Why some alpha particles were deflected

They were repelled by the positively charged nucleus

Electron configuration

Arrangement of electrons in shells (e.g. 2,8,1)

Why atoms form ions

To get a full outer shell and become more stable

Group 1 ions

+1 (lose one electron)

Group 7 ions

-1 (gain one electron)

Ionisation

Gain or loss of electrons

First ionisation energy

Energy needed to remove one electron from a gaseous atom

Why ionisation energy increases across a period

More protons increase attraction while shielding stays similar

Periodic table arrangement

Elements arranged by increasing atomic number

Same group elements

Same number of outer electrons

Same period elements

Same number of electron shells

Why group elements have similar properties

Same outer shell electrons

Metal properties

Good conductors, malleable, lose electrons to form positive ions

Non-metal properties

Poor conductors, brittle, gain electrons

Metalloid

Element with properties between metals and non-metals

Where metals are found

Left and centre of periodic table

Where non-metals are found

Right side of periodic table

Reactivity of Group 1

Increases down the group

Why Group 1 reactivity increases

More shells and shielding so outer electron is lost more easily

Group 1 reaction with water

Metal plus water forms metal hydroxide and hydrogen

Melting points of Group 1

Decrease down the group

Reactivity of Group 7

Decreases down the group

Why Group 7 reactivity decreases

More shielding makes it harder to gain an electron

Displacement reaction

A more reactive halogen replaces a less reactive one

Example of displacement

Chlorine plus potassium bromide forms potassium chloride and bromine

Boiling points of Group 7

Increase down the group

Why noble gases are unreactive

They have full outer electron shells

Boiling points of Group 0

Increase down the group

Why boiling points increase in Group 0

Stronger intermolecular forces

Why atoms get smaller across a period

Increased nuclear charge pulls electrons closer

Why atoms get bigger down a group

More electron shells increase size

Shielding

Inner electrons reduce attraction between nucleus and outer electrons

Nuclear charge

Total positive charge from protons

Why metals conduct electricity

They have delocalised electrons that move freely

Why ionic compounds have high melting points

Strong electrostatic forces between oppositely charged ions

Why sodium is more reactive than lithium

Sodium has more shells and shielding so it loses its outer electron more easily

Why chlorine is more reactive than iodine

Chlorine has less shielding so it attracts electrons more strongly

Why atomic radius decreases across a period

Increasing nuclear charge with similar shielding pulls electrons closer