We can use tools like the periodic table of elements to figure out exactly how many protons, and thus electrons, an atom has. First of all, we know that for an atom to have a neutral charge, it must have the same number of protons and electrons. If an atom loses or gains electrons, it becomes ionized, or charged. The periodic table will give us the atomic number of an element. The atomic number tells us how many protons an atom has. For example, hydrogen has an atomic number of one - which means it has one proton, and thus one electron - and actually has no neutrons.
For the Student Based on the previous description of the atom, draw a model of the hydrogen atom. The "standard" model of an atom is known as the Bohr model. Different forms of the same chemical element that differ only by the number of neutrons in their nucleus are called isotopes. Most elements have more than one naturally occurring isotope.
Many more isotopes have been produced in nuclear reactors and scientific laboratories. Isotopes usually aren't very stable, and they tend to undergo radioactive decay until something that is more stable is formed.
You may be familiar with the element uranium - it has several unstable isotopes, U being one of the most commonly known. The means that this form of uranium has neutrons and protons combined. If we looked up uranium's atomic number, and substracted that from , we could calculate the number of neutrons that isotope has. Here's another example - carbon usually occurs in the form of C carbon , that is, 6 protons and 6 neutrons, though one isotope is C, with 6 protons and 7 neutrons.
For the Student Use the periodic table and the names of the elements given below to figure out how many protons, neutrons and electrons they have.
Draw a model of an atom of the following element: silicon, magnesium, sulphur, oxygen, and helium For the Student Using the text, define the following terms: energy levels, absorption, emission, excited state, ground state, ionization, atom, element, atomic mass, atomic number, isotope.
A Optional Note on the Quantum Mechanical Nature of Atoms While the Bohr atom described above is a nice way to learn about the structure of atoms, it is not the most accurate way to model them. Although each orbital does have a precise energy, the electron is now envisioned as being smeared out in an "electron cloud" surrounding the nucleus. It is common to speak of the mean distance to the cloud as the radius of the electron's orbit. So just remember, we'll keep the words "orbit" and "orbital", though we are now using them to describe not a flat orbital plane, but a region where an electron has a probability of being.
Electrons are kept near the nucleus by the electric attraction between the nucleus and the electrons. Kept there in the same way that the nine planets stay near the Sun instead of roaming the galaxy. Unlike the solar system, where all the planets' orbits are on the same plane, electrons orbits are more three-dimensional. Each energy level on an atom has a different shape. There are mathematical equations which will tell you the probability of the electron's location within that orbit.
Let's consider the hydrogen atom, which we already drew a Bohr model of. Probable locations of the electron in the ground state of the Hydrogen atom. What you're looking at in these pictures are graphs of the probability of the electron's location. The nucleus is at the center of each of these graphs, and where the graph is lightest is where the electron is most likely to lie.
What you see here is sort of a cross section. That is, you have to imagine the picture rotated around the vertical axis. Kirchoff and Bunsen carefully measured the number and position of all the spectral lines they saw given off by a whole range of materials. These were called emission spectra , and when they had collected enough of them it was clear that each substance produced a very characteristic line spectrum that was unique.
No two substances produced exactly the same series of lines, and if two different materials were combined they collectively gave off all the lines produced by both substances.
This, thought Kirchoff and Bunsen, would be a good way of identifying substances in mixtures or in materials that needed to be analyzed. So they did. In they found a spectrum of lines that they had never seen before, and which did not correspond to any known substance, so, quite rightly, they deduced that they had found a new element, which they called cesium from the Latin word meaning "sky blue". Guess in what part of the spectrum they found the lines!
All the research on atomic structure and the hideously difficult-to-understand properties of electrons come together in the topic of "electron energy". An atom such as lithium has three electrons in various orbitals surrounding the atomic center.
These electrons can be bombarded with energy and if they absorb enough of the quanta of energy being transferred they jump about and in the most extreme case, leave the lithium atom completely. This is called ionization. Partly this difference in the amount of energy needed to dislodge different electrons away from the lithium atomic center is due to the fact that the center of the lithium atom is carrying the positive charges of three protons.
Moving a negatively charged electron away from a positively charged atomic center needs more and more energy as the amount of un-neutralized charge increases, thus;. However, the amount of energy needed to remove the first electron is a good measure of what it takes to stimulate an electron to leave its atom, and how tightly it is held there in the first place. Within the atom, as Bohr pointed out, there are different possible positions for electrons to be found as defined by the principal quantum number , usually written as " n ".
Bohr defined the energy of electrons located at these different locations of quantum state by the formula:. This is usually presented in the form of a diagram see left. If the quantum is too small the electron could not reach the next level, so it doesn't try.
If the quantum is too large the electrons would overshoot the next level, so again, it does not try. Only quanta of exactly the right size will be absorbed and used. But the amount of energy given off will be a whole number quantum. If this energy is given off as light such as happens with emission spectra then the photons rushing away from the falling electron will be of only one size and quality color.
Hence glowing sodium, or LEDs, only give off very discrete bands of light with distinct colors or bands within their spectrum. All this implies that if white light with all the possible wavelengths, colors and possible quanta of energy is shone on certain materials or substances only certain wavelengths and their quanta of energy will be absorbed by the electrons in that substance. Only a narrow band of light will have just the right quanta to move an electron to the next level, or the level above that, and so on.
That wavelength will be taken out of the spectrum of light and leave a dark band of no-light behind. Absorption spectroscopy, therefore, is the equal and opposite of emission spectroscopy. However, in both kinds, it is the absorption of quanta to move electrons, or the emission of quanta to move electrons around in the atom that is the reason why only certain wavelengths of light are affected.
Although Bohr's original picture of a quantum atom has been modified in the years since he first proposed the concept, never the less, the main principles still stand:. Electrons are to be found occupying certain volumes of space around an atomic center "nucleus" - these volumes of space are called orbitals. An electron in an orbital has a defined wavelength. The shape and location of the orbital is determined by the fact that the only stable shapes and locations are those where the electrons acting as waves can have a number of waves that are whole numbers technically these are called " standing waves ".
Standing-wave orbitals are the only ones in which the occupying electrons do not either radiate energy, or collapse. Click here to. Atomic Structure. The Nature of
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