Revising Atoms. Learning Objectives Candidates should be able to: Identify and describe protons,...

AS Chemistry

Revising Atoms

Learning Objectives

Candidates should be able to:

Identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses.

Deduce the behaviour of beams of protons, neutrons and electrons in electric fields.

Describe the distribution of mass and charges within an atom.

Deduce the number of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge).

Distinguish between isotopes on the basis of different numbers of neutrons present.

Starter activity

Democritus:Ancient Greek Philosopher-Scientist,

History of the Atom

‘a tomos’ – cannot be cut.

The problem: he was unable to provide the evidence needed to convince people that atoms really existed.

http://images.google.co.uk/imgres?imgurl=http://greece.mrdonn.org/greece06.gif&imgrefurl=http://greece.mrdonn.org/empire.html&h=450&w=629&sz=101&hl=en&start=1&tbnid=kH9RkhrDgt1a8M:&tbnh=98&tbnw=137&prev=/images?q=ancient+greek+&gbv=2&svnum=10&hl=en&sa=G

History of the Atom In 1808, an English school teacher

named John Dalton proposed that atoms could not be divided and that all atoms of a given element were exactly alike.

Dalton’s theory is considered the foundation for the modern atomic theory.

Dalton’s theory was developed with scientific basis and was accepted by others.

History of the AtomAt the end of the nineteenth century, a scientist called J.J. Thomson discovered the electron.

Thomson suggested that they could only have come from inside atoms. So Dalton's idea of the indestructible atom had to be revised.

Thomson imagined the electrons as the bits of plum in a plum pudding

History of the Atom

This implies a dense, positively charged central region containing most of the atomic mass and that the atom is mostly space.

In 1872-1937, Rutherford et al. ran experiments to determine the structure of an atom. When positively charged particles are fired into gold foil, most pass straight through while a few are violently deflected.

The Gold Foil Experiment

Rutherford expected the alpha particles to go straight through the gold foil.

Instead, some of the alpha particles were deflected, implying a central positively charged region (nucleus).

http://en.wikipedia.org/wiki/Image:Rutherford_gold_foil_experiment_results.svg

History of the Atom

In 1913, the Danish scientist Niels Bohr suggested that electrons in an atom move in set paths (energy levels) around the nucleus much like the planets orbit the sun.

Electrons can only be in certain energy levels and must gain energy to move to a higher energy level or lose energy to move to a lower energy level.

History of the Atom

In the 1920’s deBroglie & Shrodinger showed that the “solar system” model of the atom was incorrect. Instead, electrons orbit the nucleus in orbitals.

This is called quantum mechanics. We will look at this in our next lesson.

• For some time people thought atoms were the smallest particles and that they could not be broken into anything smaller.

• We now know that atoms are themselves made from even smaller and simpler particles.

• These particles are• Protons• Neutrons• Electrons

Even Smaller Particles!

J.J. Thompson – discovered presence of electrons and proposed ‘Plum Pudding’ model of the atom.

Rutherford’s ‘Gold foil’ experiment concluded that an atom's mass must be concentrated in a small positively charged nucleus and that most of the atom must be empty space. This space must contain the electrons.

Evidence for Sub-atomic particles

There are two properties of sub-atomic particles that are especially important:◦Mass◦Electrical charge

Properties of Sub-atomic Particles

Particle Charge Relative MassProtons +1 1Neutrons 0 1Electrons -1 negligible

Element atoms contain equal numbers of protons and electrons and so have no overall charge

A

B

C

–

+

Properties of Sub-atomic Particles

proton

electron

neutron

Protons, neutrons and electrons are NOT evenly distributed in atoms.

The protons and neutrons exist in a dense core called the nucleus.

Around the outside are very thinly spread electrons.

These electrons exist in layers called shells.

How Are the Particles Arranged?The Nucleusa dense core of protons and neutrons containing nearly all the mass of the atom

‘Shells’ of electronselectrons are really very very tiny so the atom is mostly empty space.

The atom of any particular element always contains the same number of protons. E.g.◦ Hydrogen atoms always contain 1 proton◦ Carbon atoms always contain 6 protons◦ Magnesium atoms always contain 12

protons

The number of protons in an atom is known as its atomic or proton number.

It is the smaller of two numbers shown in most periodic tables

Proton or Atomic Number

12

C6

Note that any element has a definite and fixed number of protons.

If we change the number of protons in an atom then this changes that atom into a different element.

Changes in the number of particles in the nucleus (protons or neutrons) is very rare. It only takes place in nuclear processes such as radioactive decay, nuclear bombs or nuclear reactors.

How Many Protons?

The mass of each atom results almost entirely from the number of protons and neutrons that are present. (Remember that electrons have a relatively tiny mass).

The sum of the number of protons and neutrons in an atom is the mass number.

Mass or Nucleon Number

Atom Protons Neutrons Mass NumberHydrogen 1 0 1

Lithium 3 4 7

Aluminium 13 14 27

Electrons are not evenly spread. The exist in energy levels known as shells. The arrangement of electrons in these shells is

often called the electron configuration.

How Are Electrons Arranged?

2nd Shell

1st Shell

3rd Shell

4th Shell

Each shell has a maximum number of electrons that it can hold.

How Many Electrons per Shell?

1st Shell: 2 electrons

2nd Shell: 8 electrons

3rd Shell: Initially 8 electrons

The maximum

Opposites attract. Protons are + and electrons are –

charged. Electrons will occupy the shells nearest

the nucleus unless these shells are already full.

Which Shells do Electrons go into?

1st Shell: Fills this first

2nd Shell: Fill this next

3rd Shell: And so on

How many electrons do the element atoms have? (This will equal the atomic number).

Keeping track of the total used, feed them into the shells working outwards until you have used them all up.

Working Out Electron Arrangements

1st Shell: Fills this first

2nd Shell: Fill this next

Drawing neat diagrams helps you keep track!

It is not strictly true to say that elements consist of one type of atom.

Whilst atoms of a given element always have the same number of protons, they may have different numbers of neutrons.

Atoms that differ in this way are called isotopes.

How Many Neutrons 1

Remember: The number of protons defines the element

Isotopes are virtually identical in their chemical reactions. (There may be slight differences in speeds of reaction).

This is because they have the same number of protons and the same number of electrons.

The uncharged neutrons make no difference to chemical properties but do affect physical properties such as melting point and density.

How Many Neutrons 2

Natural samples of elements are often a mixture of isotopes. About 1% of natural carbon is carbon-13.

Isotopes: Carbon

Protons

Electrons

Neutrons

C12

699% C

13

61%

6

6

6

6

6

7

Hydrogen exists as 3 isotopes although Hydrogen-1 makes up the vast majority of the naturally occurring element.

Isotopes: Hydrogen

H1

1

ProtonsElectronsNeutrons

Hydrogen

H2

1


(Deuterium)

H3

1


(Tritium)

About 75% of natural chlorine is 35Cl the rest is 37Cl.

Isotopes: Chlorine

Cl35

17

75%

17ProtonsElectronsNeutrons

17

18


171720

Cl37

1725%

AS Chemistry

Atomic Orbitals

Learning Objectives

Candidates should be able to:

Describe the number and relative energies of the s, p, and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals.

Describe the shapes of s and p orbitals.

State the electronic configuration of atoms given the proton number.

Starter activity

An electron’s exact location cannot be determined.

Imagine the moving blades of a fan – If you were asked where any one of the blades was located at a certain instant, you would not be able to give an exact answer – the blades are moving too quickly!

It is the same with electrons –the best a scientist can do is calculate the chance of finding an electron in a certain place within an atom

Location of Electrons – The Problem with Bohr’s Model

Energy levels and sub-levels

Energy levels

These are broadly similar to the “shells” used in GCSE Chemistry

You need to know about energy levels 1, 2, 3 and 4 at A-level

Energy level 1 is lowest in energy and closest to the nucleus


Sub-levels

The main energy levels contain sub-levels

The different main energy levels have different sub-levels in them

There are four types: s, p, d, f


Main energy level

Sub-levelsMax. no. of

electron pairs in sub-level

Max. no. of electrons in

sub-level

Max. no. of electrons in main level


Main energy level




sub-level


1


Main energy level




sub-level


1 s


Main energy level




sub-level


1 s 1


Main energy level




sub-level


1 s 1 2


Main energy level




sub-level


1 s 1 2 2


Main energy level




sub-level


1 s 1 2 2

2


Main energy level




sub-level


1 s 1 2 2

2s 1 2


Main energy level




sub-level


1 s 1 2 2

2s 1 2

p


Main energy level




sub-level


1 s 1 2 2

2s 1 2

p 3


Main energy level




sub-level


1 s 1 2 2

2s 1 2

p 3 6


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

p 3 6


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

p 3 6

d


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

p 3 6

d 5


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

p 3 6

d 5 10


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

18p 3 6

d 5 10


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

18p 3 6

d 5 10

4


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

18p 3 6

d 5 10

4

s 1 2

p 3 6

d 5 10


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

18p 3 6

d 5 10

4

s 1 2

p 3 6

d 5 10

f


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

18p 3 6

d 5 10

4

s 1 2

p 3 6

d 5 10

f 7


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

18p 3 6

d 5 10

4

s 1 2

p 3 6

d 5 10

f 7 14


Main energy level




sub-level


1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

18p 3 6

d 5 10

4

s 1 2

32p 3 6

d 5 10

f 7 14

Hydrogen's electron - the 1s orbital

Spherical

2s orbital

2p orbital

Dumb-bell shaped

p-orbitals – have direction

More complex orbitals

http://winter.group.shef.ac.uk/orbitron/

http://winter.group.shef.ac.uk/orbitron/

The order of filling

1s


1s

2s

2p


1s

2s

2p

3s

3d

3p


1s

2s

2p

3s

3d

3p4s

4p

4d

4f


1s

2s

2p

3s

3d

3p4s

4p

4d5s

5p

4f6s


1s

2s

2p

3s

3d

3p4s

4p

Electrons fill the lowest available energy level

4s fills before 3d

Electrons remain unpaired as far as possible

Cr an electron is promoted from 4s to 3d to give a half-filled 3d subshellCu an electron is promoted from 4s to 3d to give a full 3d subshell

Click to add electrons


1s

2s

2p

3s

3d

3p4s

4p

Electronic configuration in shorthand nomenclature

Click to add electrons

H 1s1He 1s2Li 1s2 2s1Be 1s2 2s2B 1s2 2s2 2p1C 1s2 2s2 2p2N 1s2 2s2 2p3O 1s2 2s2 2p4F 1s2 2s2 2p5Ne 1s2 2s2 2p6Na 1s2 2s2 2p6 3s1Mg 1s2 2s2 2p6 3s2Al 1s2 2s2 2p6 3s2 3p1 Si 1s2 2s2 2p6 3s2 3p2 P 1s2 2s2 2p6 3s2 3p3 S 1s2 2s2 2p6 3s2 3p4 Cl 1s2 2s2 2p6 3s2 3p5 Ar 1s2 2s2 2p6 3s2 3p6 K 1s2 2s2 2p6 3s2 3p6 4s1 Ca 1s2 2s2 2p6 3s2 3p6 4s2 Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1 Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2 V 1s2 2s2 2p6 3s2 3p6 4s2 3d3 Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Mn 1s2 2s2 2p6 3s2 3p6 4s2 3d5 Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7 Ni 1s2 2s2 2p6 3s2 3p6 4s2 3d8 Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10 Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10 Ga 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1Ge 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2As 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3Se 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5Kr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6


1s

2s

2p

3s

3d

3p4s

Zn Zn2+

4s electrons (outer shell) are removed before 3d (inner shell)

Ionisation

The order of filling - ionisation

AS Chemistry

Ionisation Energy

Learning Objectives

Candidates should: Be able to explain and use the term first

ionisation energy. Know the factors which effect the first

ionisation energies of elements.

Be able to explain the trend in first ionisation energies across a period and down a group of the Periodic Table.

Starter Activity

The first ionisation energy

This is the energy required to remove the outermost electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

This is more easily seen in symbol terms.

Factors affecting the size of the ionisation energy

The charge on the nucleus.

The distance of the electron from the nucleus.

The number of electrons between the outer electrons and the nucleus, i.e. the shielding.

Whether the electron is on its own in an orbital or paired with another electron (electronic repulsion).

0

200

400

600

800

1000

1200

1400

1600

Na Mg Al Si P S Cl Ar

Element

Firs

t io

nisa

tion

ene

rgy

(kJ m

ol-1

)First ionisation energies of the Group 3 elements

1st I.E INCREASES as you go across Period 3:

General Trend

there are more protons in each nucleus so the nuclear charge in each element increases ...

therefore the force of attraction between the nucleus and outer electron is increased, and ...

there is a negligible increase in shielding because each successive electron enters the same energy level ...

so more energy is needed to remove the outer electron.

Magnesium to aluminium

Look at their electronic configurations:

Magnesium: 1s2 2s2 2p6 3s2 ... and ... Aluminium: 1s2 2s2 2p6 3s2 3p1

The outer electron in aluminium is in a p sub-level. This is higher in energy than the outer electron in magnesium, which is in an s sub-level, so less energy is needed to remove it.

Look at their electronic configurations:

Phosphorus: 1s2 2s2 2p6 3s2 3p3 ... and ... Sulphur: 1s2 2s2 2p6 3s2 3p4

It's not immediately obvious what's going on until we look at the arrangements of the electrons:

Phosphorus to sulphur

Phosphorus to sulphur

The 3p electrons in phosphorus are all unpaired.

In sulphur, two of the 3p electrons are paired.

There is some repulsion between paired electrons in the same sub-level.

This reduces the force of their attraction to the nucleus, so less energy is needed to remove one of them

Al(g) Al+(g) + e- 1st I.E. = 577 kJ mol-1

Al+(g) Al2+(g) + e- 2nd I.E. = 1820 kJ mol-1

Al2+(g) Al3+

(g) + e- 3rd I.E. = 2740 kJ mol-1

Al3+(g) Al4+

(g) + e- 4th I.E. = 11600 kJ mol-1

Successive ionisation energies

You can have as many successive ionisation energies as there are electrons in the original atom.

Al(g) Al3+(g) + 3e-

If you wanted to form an Al3+(g) ion from Al(g) you

would have to supply 577 + 1820 + 2720 = +5117 kJ mol-1 of energy.

Revising Atoms. Learning Objectives Candidates should be able to: Identify and describe protons,...

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