Positively charged atoms called cations are formed when an atom loses one or more electrons. An atom that gains one or more electrons will exhibit a negative charge and is called an anion. The charge of an atom is defined as follows:Ītomic charge = number of protons − number of electronsĪs will be discussed in more detail later in this chapter, atoms (and molecules) typically acquire charge by gaining or losing electrons. When the numbers of these subatomic particles are not equal, the atom is electrically charged and is called an ion. This “missing” mass is known as the mass defect, and you will learn about it in the chapter on nuclear chemistry.)Ītomic number ( Z ) = number of protons mass number ( A ) = number of protons + number of neutrons A − Z = number of neutrons atomic number ( Z ) = number of protons mass number ( A ) = number of protons + number of neutrons A − Z = number of neutronsĪtoms are electrically neutral if they contain the same number of positively charged protons and negatively charged electrons. (An observant student might notice that the sum of an atom’s subatomic particles does not equal the atom’s actual mass: The total mass of six protons, six neutrons, and six electrons is 12.0993 amu, slightly larger than 12.00 amu. The properties of these fundamental particles are summarized in Table 2.2. The electron has a charge of 1− and is a much lighter particle with a mass of about 0.00055 amu (it would take about 1800 electrons to equal the mass of one proton). A neutron is a slightly heavier particle with a mass 1.0087 amu and a charge of zero as its name suggests, it is neutral. (The Dalton (Da) and the unified atomic mass unit (u) are alternative units that are equivalent to the amu.) The fundamental unit of charge (also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 × × 10 −19 C.Ī proton has a mass of 1.0073 amu and a charge of 1+. (This isotope is known as “carbon-12” as will be discussed later in this module.) Thus, one amu is exactly 1 12 1 12 of the mass of one carbon-12 atom: 1 amu = 1.6605 × × 10 −24 g. Since 1961, it has been defined with regard to the most abundant isotope of carbon, atoms of which are assigned masses of exactly 12 amu. The amu was originally defined based on hydrogen, the lightest element, then later in terms of oxygen.
When describing the properties of tiny objects such as atoms, we use appropriately small units of measure, such as the atomic mass unit (amu) and the fundamental unit of charge (e). For example, a carbon atom weighs less than 2 × × 10 −23 g, and an electron has a charge of less than 2 × × 10 −19 C (coulomb). (credit middle: modification of work by “babyknight”/Wikimedia Commons credit right: modification of work by Paxson Woelber)Ītoms-and the protons, neutrons, and electrons that compose them-are extremely small. For a perspective about their relative sizes, consider this: If the nucleus were the size of a blueberry, the atom would be about the size of a football stadium ( Figure 2.11).įigure 2.11 If an atom could be expanded to the size of a football stadium, the nucleus would be the size of a single blueberry.
The diameter of an atom is on the order of 10 −10 m, whereas the diameter of the nucleus is roughly 10 −15 m-about 100,000 times smaller. The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. It was learned that an atom contains a very small nucleus composed of positively charged protons and uncharged neutrons, surrounded by a much larger volume of space containing negatively charged electrons. The development of modern atomic theory revealed much about the inner structure of atoms.
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