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.

 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 Atom  At 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

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

 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).

 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.

 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

 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.

 There are two properties of sub-atomic particles that are especially important: ◦ Mass ◦ Electrical chargeParticleCharge Relative Mass Protons+11 Neutrons01 Electronsnegligible Element atoms contain equal numbers of protons and electrons and so have no overall charge

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. The Nucleus a dense core of protons and neutrons containing nearly all the mass of the atom ‘Shells’ of electrons electrons are really very very tiny so the atom is mostly empty space.

always  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 12 C 6

 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. is very rare  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.

 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  The sum of the number of protons and neutrons in an atom is the mass number. AtomProtonsNeutronsMass Number Hydrogen101 Lithium347 Aluminium131427

 Electrons are not evenly spread.  The exist in energy levels known as shells. electron configuration.  The arrangement of electrons in these shells is often called the electron configuration. 2nd Shell 1st Shell 3rd Shell 4th Shell

 Each shell has a maximum number of electrons that it can hold. 1st Shell: 2 electrons 2 nd Shell: 8 electrons 3 rd 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. 1st Shell: Fills this first 2 nd Shell: Fill this next 3 rd 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. 1st Shell: Fills this first 2 nd 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. isotopes.  Atoms that differ in this way are called isotopes. 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.

 Natural samples of elements are often a mixture of isotopes. About 1% of natural carbon is carbon-13. Protons Electrons Neutrons C % C %

Hydrogen exists as 3 isotopes although Hydrogen-1 makes up the vast majority of the naturally occurring element. H 1 1 Protons Electrons Neutrons Hydrogen H 2 1 Protons Electrons Neutrons (Deuterium) H 3 1 Protons Electrons Neutrons (Tritium)

About 75% of natural chlorine is 35 Cl the rest is 37 Cl. Cl % 17 Protons Electrons Neutrons Protons Electrons Neutrons Cl %

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 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-levels Max. 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-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s1

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s12

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 p

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 p3

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 p36

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 p36

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 p36 d

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 p36 d5

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 p36 d510

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 18 p36 d510

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 18 p36 d510 4

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 18 p36 d510 4 s12 p36 d5

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 18 p36 d510 4 s12 p36 d5 f

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 18 p36 d510 4 s12 p36 d5 f7

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 18 p36 d510 4 s12 p36 d5 f714

Main energy level Sub-levels Max. no. of electron pairs in sub-level Max. no. of electrons in sub-level Max. no. of electrons in main level 1s122 2 s12 8 p36 3 s12 18 p36 d510 4 s12 32 p36 d510 f714

Hydrogen's electron - the 1s orbital Spherical

2s orbital

Dumb-bell shaped

More complex orbitals

The order of filling

1s The order of filling

1s 2s 2p The order of filling

1s 2s 2p 3s 3d 3p The order of filling

1s 2s 2p 3s 3d 3p 4s 4p 4d 4f The order of filling

1s 2s 2p 3s 3d 3p 4s 4p 4d 5s 5p 4f 6s The order of filling

1s 2s 2p 3s 3d 3p 4s 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 subshell Cu an electron is promoted from 4s to 3d to give a full 3d subshell Click to add electrons The order of filling

1s 2s 2p 3s 3d 3p 4s 4p Electronic configuration in shorthand nomenclature Click to add electrons H 1s 1 He 1s 2 Li 1s 2 2s 1 Be 1s 2 2s 2 B 1s 2 2s 2 2p 1 C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6 Na 1s 2 2s 2 2p 6 3s 1 Mg 1s 2 2s 2 2p 6 3s 2 Al 1s 2 2s 2 2p 6 3s 2 3p 1 Si 1s 2 2s 2 2p 6 3s 2 3p 2 P 1s 2 2s 2 2p 6 3s 2 3p 3 S 1s 2 2s 2 2p 6 3s 2 3p 4 Cl 1s 2 2s 2 2p 6 3s 2 3p 5 Ar 1s 2 2s 2 2p 6 3s 2 3p 6 K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Sc 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 Ti 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 V 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 Cr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Mn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Co 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 Ni 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Zn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 Ga 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 Ge 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 As 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Se 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 Br 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 Kr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 The order of filling

1s 2s 2p 3s 3d 3p 4s ZnZn 2+ 4s electrons (outer shell) are removed before 3d (inner shell) Ionisation The order of filling - ionisation

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.

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).

First ionisation energies of the Group 3 elements

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.

Look at their electronic configurations: Magnesium: 1s 2 2s 2 2p 6 3s 2... and... Aluminium: 1s 2 2s 2 2p 6 3s 2 3p 1 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: 1s 2 2s 2 2p 6 3s 2 3p 3... and... Sulphur: 1s 2 2s 2 2p 6 3s 2 3p 4 It's not immediately obvious what's going on until we look at the arrangements of the electrons:

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)  Al 2+ (g) + e - 2nd I.E. = 1820 kJ mol -1 Al 2+ (g)  Al 3+ (g) + e - 3rd I.E. = 2740 kJ mol -1 Al 3+ (g)  Al 4+ (g) + e - 4th I.E. = kJ mol -1 Successive ionisation energies You can have as many successive ionisation energies as there are electrons in the original atom. Al (g)  Al 3+ (g) + 3e - If you wanted to form an Al 3+ (g) ion from Al (g) you would have to supply = kJ mol -1 of energy.