Atomic theory
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Atomic Theory. Everything written in black has to go into your notebook Everything written in blue should already be in there. Atom : Smallest part of any element, which can take part in a chemical reaction Element : Pure substance consisting of one type of atom only

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Atomic Theory

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Atomic Theory

Everything written in black has to go into your notebook

Everything written in blue should already be in there


  • Atom: Smallest part of any element, which can take part in a chemical reaction

  • Element: Pure substance consisting of one type of atom only

  • Compound: Substance consisting of two or more elements (e.g. H2O)

  • Mixture: Consists of two or more substances but they are not joined together chemically (e.g. sand and iron filings)


Molecule: Smallest part of a substance which can exist on its own (e.g. Cu, H2, H2O)

  • There are monatomic, diatomic or triatomic molecules


1 a.m.u. = 1.66 x 10-27 kg


The Bohr Model

Protons and neutrons are located in the nucleus of the atom

Electrons orbit the nucleus in ‘shells’ or ‘energy levels’


1st energy level can hold 2 electrons

2nd energy level can hold 8 electrons

3rd energy level can hold 18 electrons


7

3

3P 4N

E.g. Lithium Li


23

11

11P 12N

E.g. sodium Na


  • Isotopes: Atoms with the same number of protons but different numbers of neutrons (e.g. carbon 12, carbon 13 and carbon 14)


  • Isotopes: Atoms having the same number of protons but different numbers of neutrons (e.g. carbon 12, carbon 13 and carbon 14)

  • Atomic Number (Z): The number of protons in an atom

  • Mass Number (A): The number of protons and neutrons in an atom

  • Relative atomic mass: The mass of an atom compared to of the mass of an atom of carbon


  • An energy level is a specific level of energy which an electron has in an atom


The emission spectrum is the series of lines which are given out by excited atoms of an element


E2

Higher energy level

Photon of light emitted

Lower energy level

E1


E2 – E1 = hf

E2 = energy of electron in higher energy level

E1 = energy of electron in lower energy level

h = Planck’s constant

f = frequency of light emitted


Flame tests (page 126)


The ground state is the energy level of the electron before it gains energy

The excited state is the energy level of the electron after it gains energy


The Aufbau Principal

  • Electrons always fill the lowest energy level available when the atom is in the ground state

  • Remember that lower energy levels are nearer the nucleus, and higher energy levels are further away from the nucleus


2

8

18

  • 1st shell (n =1) holds up to electrons

  • 2nd shell (n=2) holds up electrons

  • 3rd shell (n=3) holds up to electrons

  • 4th shell (n=4) holds up to electrons

  • Each of these main energy levels (shells) contains “sub-levels”

32


Consider it like 5th year is split into 3 classes; so an energy level is split into ‘sub-energy levels’


4f

4d

N=4

4p

s holds 2

p holds 6

d holds 10

f holds 14

4s

3d

N=3

3p

3s

2p

N=2

2s

N=1

1s


They are filled in this order

4f

4d

4p

3d

3d “jumps” up!

4s

3p

3s

2p

2s

1s


Writing electronic configurations:

Example 2

Cobalt (Co)

1s2 2s2 2p6 3s2 3p6 4s2 3d7


  • Copper and chromium are the only exceptions to the trend


  • 1s2, 2s2, 2p6, 3s2, 3p6, 4s2,

4s1, 3d5

3d4


  • 1s2, 2s2, 2p6, 3s2, 3p6, 4s2,

4s1, 3d10

3d9


Atomic Orbitals

  • An atomic orbital is the region around a nucleus where there is a high probability of finding an electron


Main energy level

Atomic orbitals

Sub-level

s

s

px

py

pz


s

s

px

py

pz


s

s

px

py

pz

s

px

py

pz


  • Every orbital can hold a maximum of 2 electrons


s sub-level has 1 orbital

  • spherical in shape


p sub-level has 3 orbitals

  • dumb bell shape

px py pz


d sub-level has 5 orbitals


Hund’s Rule of Maximum Multiplicity

  • When two or more orbitals of equal energy are available to electrons the electrons occupy them singly, before filling them in pairs

Sometimes called the“Bus Seat Rule”


Pauli’s Exclusion Principle

  • No more than two electrons can occupy an orbital, and they can only do this if they have opposite spin


Example

  • Oxygen

  • O


1s2 2s2 2p4


Quantum Numbers

  • A code consisting of 4 numbers, which give the full information about any one electron in an atom


Question

  • Give the 4 quantum numbers for each of the electrons in the Berylium atom


electronic configuration: 1s2 2s2

1st electron: 1, 0, 0, ½

2nd electron: 1, 0, 0, -½

3rd electron: 2, 0, 0, ½

4th electron: 2, 0, 0, -½


The Periodic Table

  • Groups go from top to bottom

  • Elements in the same group have the same number of electrons in their outer shell, and so have similar chemical properties

  • NB: Group 1 (Alkali metals)

  • Group 2 (Alkaline earth metals)

  • Group 7 (Halogens)

  • Group 8 (Noble gases)


  • Periods go from left to right

  • Elements in the same period have the same number of electron shells occupied


A group is a vertical column in the Periodic Table

A period is a horizontal row in the Periodic Table


Valency

The number of electrons that one atom of that element will give, take or share when it forms a chemical bond


Metals and Non-Metals

  • The periodic table may be divided up into metals and non-metals

  • Metals are hard, shiny, and good conductors of electricity

  • Carbon exhibits properties of both metals and non-metals


Transition Elements (d-block elements)

  • They are all metals

  • They act as catalysts

  • They form coloured compounds

  • They have an incomplete d subshell

  • They have variable valency

    • E.g. Copper, 1 or 2

    • Iron, 2 or 3


Ionisation Energy

  • The ionisation energy of an element is the minimum energy required to remove the outermost electron from a neutral gaseous atom

  • The ionisation energy is the minimum energy required to remove the outermost electron from an atom

HL

OL


  • The noble gases have very high values because they have a full outer shell and so it is difficult to remove an electron

  • The alkali metals have low values because they have one electron in their outer shell, so it is easily removed


  • Ionisation energies decrease as you go down through a group because

    • Atomic radius increases

    • The full inner shells have a screening effect


  • Ionisation energies increase as you go across a period

    • The atomic radius decreases, so it is more difficult to remove an electron

    • Exceptions are Be, N, Mg, P


Full s sublevel

  • Be: 1s2, 2s2

  • Mg: 1s2, 2s2, 2p6, 3s2

  • N: 1s2, 2s2, 2p3

  • P: 1s2, 2s2, 2p6, 3s2, 3p3

  • Their outer sub-energy levels are full so they are stable

Half-full p sublevel


Second Ionisation Energy

  • The second ionisation energy of an element is the minimum energy required to remove the second outermost electron from a neutral gaseous atom

  • The second ionisation energy is the minimum energy required to remove the second outermost electron from an atom

HL

OL


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