Gram atomic mass unit

The formula weight of a substance is equal to its number of grams per mole. Avogadro s number is the number of atomic mass units in 1 g. It is defined in that manner so that the atomic weight of an element (in amu) is numerically equal to the number of grams of the element per mole. Consider helium, with atomic weight 4.0 ... 

Avogadro s number a mole 6.02 X 1023 units the number of atomic mass units per gram. 

Our modern model describes the atom as an electrically neutral sphere with a tiny nucleus in the center containing positively charged protons and neutral neutrons. The negatively charged electrons are moving in complex paths outside the nucleus in energy levels at different distances from the nucleus. These subatomic particles have very little mass expressed in grams so we often use the unit of an atomic mass unit (amu or simply u). An amu is 1/12 the mass of a carbon atom that contains six protons and six neutrons. Table 2.1 summarizes the properties of the three subatomic particles. 

The substance s molar mass is the mass in grams of the substance that contains one mole of that substance. In the previous chapter, we described the atomic mass of an element in terms of atomic mass units (amu). This was the mass associated with an individual atom. At the microscopic level, we can calculate the mass of a compound by simply adding together the masses in amu s of the individual elements in the compound. However, at the macroscopic level, we use the unit of grams to represent the quantity of a mole. 

The mole (mol) is the amount of a substance that contains the same number of particles as atoms in exactly 12 grams of carbon-12. This number of particles (atoms or molecules or ions) per mole is called Avogadro s number and is numerically equal to 6.022 x 1023 particles. The mole is simply a term that represents a certain number of particles, like a dozen or a pair. That relates moles to the microscopic world, but what about the macroscopic world The mole also represents a certain mass of a chemical substance. That mass is the substance s atomic or molecular mass expressed in grams. In Chapter 5, the Basics chapter, we described the atomic mass of an element in terms of atomic mass units (amu). This was the mass associated with an individual atom. Then we described how one could calculate the mass of a compound by simply adding together the masses, in amu, of the individual elements in the compound. This is still the case, but at the macroscopic level the unit of grams is used to represent the quantity of a mole. Thus, the following relationships apply ... 

The total mass of an atom is called its atomic mass. This is the sum of the masses of all the atom s components (electrons, protons, and neutrons). Because electrons are so much less massive than protons and neutrons, their contribution to atomic mass is negligible. As we explore further in Section 9.2, a special unit has been developed for atomic masses. This is the atomic mass unit, amu, where 1 atomic mass unit is equal to 1.661 X 10-24 gram, which is slightly less than the mass of a single proton. As shown in Figure 3.21, the atomic masses listed in the periodic table are in atomic mass units. As is explored in the Calculation Corner on page 95, the atomic mass of an element as presented in the periodic table is actually the average atomic mass of its various isotopes. 

Was this your answer Both terms include the word mass and so are easily confused. Focus your attention on the second word of each term, however, and you ll get it right every time. Mass number is a count of the number of nucleons in an isotope. An atom s mass number requires no units because it is simply a count. Atomic mass is a measure of the total mass of an atom, which is given in atomic mass units. If necessary, atomic mass units can be converted to grams using the relationship i atomic mass unit = 1.661 x io 24 gram. 

As Figure 9.4 illustrates, if you express the numeric value of the atomic mass of any element in grams, the number of atoms in a sample of the element having this mass is always 6.02 X 1023, which is 1 mole. For example, a 22.990-gram sample of sodium metal, Na (atomic mass 22.990 atomic mass units), contains 6.02 X 1023 sodium atoms, and a 207.2-gram sample of lead, Pb (atomic mass = 207.2 atomic mass units), contains 6.02 X 1023 lead atoms. 

How many atoms are there in a 6.941-gram sample of lithium, Li (atomic mass 6.941 atomic mass units) ... 

How many molecules are there in an i8.ois-gram sample of water, H20 (formula mass 18.015 atomic mass units) ... 

What mass of water is produced when 16 grams of methane, CH4 (formula mass 16 atomic mass units), burn in the reaction... 

Two atomic mass units equal how many grams ... 

Which is greater 1.01 atomic mass units of hydrogen or 1.01 grams of hydrogen ... 

How many molecules of aspirin (formula mass 180 atomic mass units) are there in a 0.250-gram sample ... 

What mass of oxygen (in grams) is produced when 122.55 grams of KC103 (formula mass 122.55 atomic mass units) take part in this reaction ... 

If the atomic masses in the G matrix elements are expressed in atomic mass units rather than in grams and the frequencies are expressed in reciprocal centimeters (cm"1), 10.5-2 may be written... 

Atoms are so tiny that even the smallest speck of dust visible to the naked eye contains about 1016 atoms. Thus, the mass of a single atom in grams is much too small a number for convenience, and chemists therefore use a unit called an atomic mass unit (amu), also known as a dalton (Da). One amu is defined as exactly one-twelfth the mass of an atom of and is equal to 1.660 539 X 10 24 g ... 

About 1837 electrons are equal in mass to the mass of one proton or one neutron. A summary of each type of particle, its mass and relative charge is shown in Table 3.1. You will notice that the masses of all these particles are measured in atomic mass units (amu). This is because they are so light that their masses cannot be measured usefully in grams. 

Proton—a positively charged particle located in the atom s nucleus. The electrical charge has a magnitude of+ 1.6 X 10 19 coulombs (C) however, for simplicity, it is often referred to by its relative charge of +1.0 (charge relative to an electron). The mass of a proton is about 1.67 X 10 24 g. The gram is not a practical unit to describe the mass of subatomic particles, so instead we use the atomic mass unit, or amu. An amu is defined as the mass of a carbon atom containing 6 protons and 6 neutrons. The mass of a proton is 1.0073 amu. 

It is inconvenient to measure the mass of subatomic particles using units such as grams. Instead, chemists use a unit called an atomic mass unit (symbol u). A proton has a mass of about 1 u, which is equal to 1.66 x 10 24 g. 

You would never express the mass of a lump of gold, like the one in Figure 5.11, in atomic mass units. You would express its mass in grams. How does the mole relate the number of atoms to measurable quantities of a substance The definition of the mole pertains to relative atomic mass, as you learned in section 5.1. One atom of carbon-12 has a mass of exactly 12 u. Also, by definition, one mole of carbon-12 atoms (6.02 x 1023 carbon-12 atoms) has a mass of exactly 12 g. 

The Avogadro constant is the factor that converts the relative mass of individual atoms or molecules, expressed in atomic mass units, to mole quantities, expressed in grams. 

How can you use this relationship to relate mass and moles The periodic table tells us the average mass of a single atom in atomic mass units (u). For example, zinc has an average atomic mass of 65.39 u. One mole of an element has a mass expressed in grams numerically equivalent to the element s average atomic mass expressed in atomic mass units. One mole of zinc atoms has a mass of 65.39 g. This relationship allows chemists to use a balance to count atoms. You can use the periodic table to determine the mass of one mole of an element. 

The molecular mass of a compound is measured in atomic mass units but its molar mass is measured in grams. Explain why this... 

Atomic mass The at-rest mass of an atom. It is usually measured in atomic mass units or amu, which is defined as exactly one-twelfth the mass of an atom of carbon-12, the isotope of carbon with six protons and six neutrons in its nucleus. One amu is equal to approximately 1.66 x 10 24 grams. 

It is also possible to think of a mole as the number of atomic mass units in 1 gram. 

Use the value of Avogadro s number, the mass of one C atom (exactly 12 amu), and the definition of a mole to calculate the number of atomic mass units per gram. 

---