Chapter 2, Section 3. Atomic Structure and Symbolism. Pause to read the learning objectives. Atoms are very small, with a diameter in the 10 to the negative 10th meter scale. The nucleus is about a hundred thousand times smaller than that, at approximately 10 to the negative 15th meters. This means that if an atom could be expanded to the size of a football stadium, the nucleus would be the size of a single blueberry.
The rest of the atom is basically empty space. Most of the mass of the atom is in the nucleus, where protons and neutrons are located. Electrons move around the empty space around the nucleus and are much lighter than the other subatomic particles.
We need small units to refer to atoms and subatomic particles. The fundamental unit of charge, E, is 1.602 times 10 to the negative 19th coulombs. very close to what Millikan calculated. We can then say that the charge of an electron is negative one of a proton, positive one, and neutrons have no charge.
For mass, we use atomic mass unit, which is based on 1 12 of a carbon 12 atom. A carbon atom weighs less than 2 times 10 to the negative 23rd grams. That equates to 1 amu being 1.6605 times 10 to the negative 24th grams. Electrons are very lightweight at 0.00055 amu.
Protons are over 1800 times heavier at 1.0073 amu. Neutrons are slightly heavier than protons, weighing 1.0083 amu. The number of protons in the atom is its most important trait.
All atoms of a given chemical element have the same number of protons. We define as atomic number, Z, the number of protons in the atom. For example, it is known that the atomic number of carbon is 6, so any atom that contains 6 protons is the element carbon, even though the number of neutrons and or electrons may vary.
An atom with seven protons would instead be a nitrogen atom. Because the mass of an electron is so much less than that of nuclear subatomic particles, the mass number A is by convention given by the total number of protons and neutrons in an atom. Therefore, if we know the atomic number and the mass number, we can calculate the number of neutrons in that atom.
The atomic number is equal the number of protons, the mass number is equal to the number of protons plus the neutrons, therefore the number of neutrons is the mass number minus the atomic number. Here's an example. If we know an atom has six protons and a mass number of 12, we can deduce that the number of neutrons will be 12 minus 6, which is 6. Atoms that have the same number of positive and negative particles are in neutral state. They have the exact same number of protons and electrons. In this case, the atomic number, which represents the number of protons, also indicates the number of electrons in the atom.
If a neutral atom has an atomic number of 6, a mass number of 12, the number of protons is equal to 6, as is the number of electrons. The mass number does not affect the number of electrons at all. And it's just is extra information here. Atoms and molecules can acquire charge by losing or gaining electrons. When the number of protons and electrons are not equal, the atom is called an ion.
We can calculate the charge of ions by subtracting the number of electrons from the number of protons. If an atom loses an electron, it is called a cation. Since now there are more protons than electrons, the overall charge of the ion is positive.
Remember that electrons are negative, so losing negative charge yields positive charge. Here's an example. Suppose an atom of sodium, Na, that has an atomic number 11 loses one electron.
It has 11 protons but only 10 electrons. That means that the charge of the ion is 11 minus 10 or plus one, a cation. On the other hand, if an ion is formed From an atom that gained an electron, it is called anion. This ion has more electrons than protons. which gives it a negative charge.
An atom of oxygen, O, that has an atomic number of 8, gains 2 electrons. The number of protons is always the atomic number, so it is 8. The neutral atom has 8 electrons, but having gained 2, it has now 10 electrons. The charge of this ion is therefore negative 2. A chemical symbol is an abbreviation that we use to indicate an element or an atom of an element.
Here you can see some common elements and their symbols. Some symbols are derived from the common name of the element. For example, aluminum is AL. Others are abbreviations of the name in another language, such as PB for lead. Most symbols have one or two letters.
The three-letter symbols have been used to describe recently synthesized elements with atomic number greater than 1 12th. Only the first letter is capitalized in a chemical symbol. We represent an atom of a given element by the symbol of that element, the mass number as the left superscript, right here, and the atomic number as a left subscript, which is sometimes omitted. The charge is shown as a right superscript. When no charge is shown, the atom is neutral.
Here we see the element helium, which has an atomic number of two, a mass number of four, and a charge of positive two, a cation. Many elements have atoms with different mass number isotopes. Here we see three isotopes of hydrogen. All of them naturally have the exact same atomic number, as you can see. here and in the chemical symbol right here, which is one.
The number of neutrons vary though, and consequently the mass number, the mass in amu. Each of these are found in nature at different proportions, as the last column shows. Since most of the mass of the atom comes from protons and neutrons, and each has a mass of approximately 1 amu, The atomic mass of a single atom is approximately equal to its mass number.
As most elements exist naturally as a mixture of two or more isotopes, the atomic mass of an element is given by the weighted average mass of all isotopes present in naturally occurring sample of that element. This is what you find in the periodic table. This equation shows us how to calculate the atomic mass from information of the isotopes.
We sum the contributions of each isotope as the product of its fractional abundance, or how much of it we find in nature, and the mass of that isotope in AMU. As an example, let's calculate the mass of boron, which is composed of these two isotopes. The contribution of B10 is 0.199, which is 19.9% divided by 100. We multiply that by the mass number.
which is 10.0129 amu. To that, we will add the contribution of the isotope B11, which is 0.801 times 11.0093 amu. Computing this, we find that the atomic mass of boron is 10.81 amu. The occurrence and natural abundances of isotopes can be experimentally determined using an instrument called a mass spectrometer, MS. A vaporized sample is exposed to a high-energy electron beam that causes the sample's atoms, or molecules, to become electrically charged, typically by losing one or more electrons. These cations are then separated by their mass and charge.
A computer software generates a graph called a mass spectrum, which here shows the analysis for zirconium with peaks for each different isotope of Zr.