25.2 Beams of charged particles
Candidates should be able to:
(a) show an understanding of the main principles of determination of e by Millikan’s experiment
(b) summarise and interpret the experimental evidence for quantisation of charge
(c) describe and analyse qualitatively the deflection of beams of charged particles by uniform electric and uniform magnetic fields
(d) explain how electric and magnetic fields can be used in velocity selection
(e) explain the main principles of one method for the determination of v and e /me for electrons
So what happens when the photons arrive? Let‘s look back at the case of the zinc plate at the start of this section. The photons arrive and interact with (hit !) electrons in the metal. Now here‘s a useful rule to follow.
Each photon only interacts with one electron.
It delivers its energy to the electron and disappears (because it is a packet of pure energy and nothing else.) So the electron now has extra energy. What can it do with it? Well, if it has enough extra energy it can leave the metal atom. Of course that means that only photons delivering enough energy will cause electrons to leave the metal.
So only photons above the threshold frequency, fo, will cause photoelectric emission.
How come more intense (brighter) radiation doesn‘t cause emission?
More intense radiation simply means that more packets of energy (photons) are delivered each second. But the energy of each packet is unchanged. So if there wasn‘t enough energy to cause photoelectric emission, making it brighter won‘t change anything.
Assuming we‘re above the threshold frequency, there are other things to think about when the photon arrives.
Photons can easily pass between or even through atoms without hitting (interacting with) an electron. So when they do finally hit an electron, there are a number of possible scenarios.
1. The photon hits an electron at the surface of the metal. The electron uses the energy that it has gained to leave the atom and head off to freedom. That‘s an easy one.
2. The photon hits an electron at the surface of the metal. The electron leaves the atom but heads off deeper into the metal and never manages to escape.
3. The photon passes deep into the metal before it hits an electron. The electron leaves the atom and heads towards the surface and escapes!
4. The photon passes deep into the metal before it hits an electron. The electron leaves the atom and heads towards the surface but it doesn‘t have enough energy to push its way past all the other atoms to get to the surface so it grinds to a halt still inside the metal and never escapes.
5. The photon passes deep into the metal before it hits an electron. The electron leaves the atom and heads off in the wrong direction and never escapes.
So in fact a very small proportion of the photons that arrive at the metal will cause photoelectrons to be emitted.
The protons and neutrons in each atom are tightly packed in a positively charged nucleus. Negatively charged electrons move around the nucleus. The number of protons in a nucleus defines the type of atom and is the same asthe atomic number. The number of neutrons is found by subtracting the atomic number from the mass number. In an atom because there is no overall charge the number of electrons equals the number of protons.
In chemical reactions the nucleus remains unchanged.
Geiger and Marsden bombarded a thin gold foil with a beam of alpha particles.
Most of the particles passed through the foil without deflection and were detected by a flash of light when the alpha particle struck a zinc sulphide screen, surrounding the gold foil.
A few were deflected and some of these were deflected at angles greater than 900, suggesting they had been repelled by large positive charges within the foil - nuclei of atoms of gold.
From GCSE you should be familiar with the Bohr model of electrons arranged around a nucleus. The electrons are in certainenergy levels and each energy level can hold only up to a maximum number of electrons.
This is summarised in the table below:
Energy level or ‘shell‘
Max no of electrons
However, these models of electron arrangement are simple and a more advance done can now be used. It is possible to break these energy levels into sub-shells.
Electrons are impossible to locate exactly at any one time. It is however, possible to indicate a region or volume where the electron is most likely to be found. This region is called an Orbital.
Each orbital is capable of holding a maximum of 2 electrons. Orbitals can be divided into s, p, d, and f types. Each type has its own characteristic shape.
The shape of s and p orbitals are shown below:
The first energy level holds a maximum of 2 electrons in one s type orbital called 1s. There are no p, d, or f orbitals available at this energy level.
The second energy level consists of one s type orbital and three p type orbitals: 2s, 2px, 2py, 2pz.
Note: there are 3 p orbitals of identical energy, one along the x axis, one along the y-axis and one along the z-axis.
These four orbitals can hold a total of 8 electrons (i.e. 2 electrons each). There are no 2d or 2f orbitals.
The third energy level consists of: one s type orbital, three p type orbitals and 5 d type orbitals. These nine orbitals can hold a maximum of 18 electrons altogether (i.e. two electrons each).
Note: there are seven f type orbitals holding a maximum of 14 electrons in total.
When filling the available orbitals with electrons two important principles should be followed:
1. Electrons fill the lowest energy orbitals first and the other orbitals in order of ascending energy. It is incorrect to assume that an energy level is always completely filled before electrons enter the next energy level. The order of filling orbitals as shown below is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p.
2. Where there are several orbitals of exactly the same energy e.g. three 2p orbitals, electrons will occupy different orbitals whenever possible.
For example: nitrogen is 1s2 2s2 2px1 2py1 2pz1 and not 1s2 2s2 2px2 2py1.
This principle is Hund’s rule. When an orbital only contains 1 electron then this electron is said to be unpaired.
a) The small number above the orbital refers to the number of electrons in the orbital: 1s2 means 2 electrons in a 1s orbital.
b) The electron arrangements are sometimes abbreviated.
For example: the electron arrangement for calcium may be written as 1s2 2s2 2p6 3s2 3p6 4s2.
Most elements form compounds.
For example: A reaction between sodium and chlorine gives the compound sodium chloride(salt) quite readily.
The noble gases do not usually form compounds. They are different from other elements, since their atoms are described as stable or unreactive. They are stable because their outer electron shell is full. A full outer shell makes an atom more stable.
Only the noble gases have full outer shells. This is why they are stable.
Other elements react with each other in order to obtain full outer shells, this makes them more stable.
Depending on their electronic configurations, atoms lose or gain electrons in order to achieve a full outer shell.
The sodium atom has one electron in its outer shell. If it loses this one electron it will achieve a full outer shell. By losing the one electron to another atom, it becomes a sodium ion.
The sodium ion still has 11 protons but by losing one electron it has only 10 electrons compared to the atom. Hence, itsoverall charge is +1.
This +1 charge is due to the ion having one more proton than electron.
In naming ions, you take the symbol Na and assign a positive charge. This gives us the sodium ion Na+.
A chlorine atom has seven electrons in its outer shell. It can reach a full outer shell by gaining one electron. It will then become the chloride ion, Cl-.
A negative charge is assigned to the ion to signify that the ion contains one more electron than proton.
Any atom can become an ion if it gains or loses electrons.
An ion is a charged particle. It is charged due to an unequal number of electrons and protons.