In chemistry, Avogadro's number is a constant numerical value that shows the number of units (like, particles, molecules, etc) of a chemical solution or a chemical element on its one-unit mole. The number of electrons in a neutral atom of an element is always equal to the of the element. A) mass number b) atomic number c) atomic mass unit d) isotope number e) Avogadro's number. What is the atomic weight of a hypothetical element consisting of two isotopes, one with mass = 64.23 amu (26.00%), and one with mass = 65.32 amu. . The mass of an atom in amu is numerically the same as the mass of one mole of atoms of the element in grams. One atom of sulfur has a mass of 32.07 amu; one mole of S atoms has a mass of 32.07 g. Mass of 1 H atom: 1.008 amu x 1.661 x10-24 g/amu = 1.674 x10-24 g Mass of 1 mole of H atoms: 1.674 x10-24g/H atom x 6.022 x1023H atoms = 1.008. Thanx for A2A As, 1 amu = 1/12 of mass of one C-12 atom The value of amu doesn't depend upon Avogadro number. So atomic mass of oxygen will still be same as 16 amu, which simply means that one atom of O is 16 times heavier than “ 1/12 of mass o.

### = 1 Gram

Molecules and atoms are extremely small objects - both in size and mass. Consequently, working with them in the laboratory requires a large collection of them. How large does this collection need to be? A standard needs to be introduced. This standard is the 'mole'. The mole is based upon the carbon-12 isotope. We ask the following question: How many carbon-12 atoms are needed to have a mass of exactly 12 g. That number is NA - Avogadro's number. Thus, NA is defined by

NA x (mass of carbon-12 atom) = 12 g

Careful measurements yield a value for NA = 6.0221367x10^+23. This is an incredibly large number - almost a trillion trillion. For example, if we stack NA pennies on top of one another how tall would the stack be? The answer is it would be so tall that the stack of pennies could reach the sun and back almost 500 million times!

A convenient name is given when there is an Avogadro's number of objects - it is called a 'mole'. Thus in the above example there was a mole of pennies.

1 mole = NA objects

The mole concept is no more complicated than the more familiar concept of a dozen : 1 dozen = 12 objects. From the penny example above one might suspect that the mass of a mole of objects is huge. Well, that is true if we're considering a mole of pennies, however a mole of atoms or molecules is a different story. Recall that the atomic mass unit (amu) is defined as 1/12 the mass of a carbon-12 atom. Consequently we have the relation

NA x 12 amu = 12 g

Thus, a mole of carbon-12 atoms has a mass of just 12 g. What about other atoms? In the periodic table the atomic mass of the elements is given. For example the atomic mass of magnesium is 24.305 amu. This is the average isotopic mass of naturally occurring magnesium. What is the molar mass of magnesium in grams? From the equation above we get 1 amu = 1g/NA or 1 amu = 1.66054x10^-24 g. Thus, a mole of magnesium atoms has a mass of NA x 24.305 amu x (1.66054x10^-24 g/amu) = 24.305 g. A mole of magnesium atoms has a mass of 24.305 g. This example demonstrates that the atomic mass of magnesium can be interpreted in one of two ways: (1) the average mass of a single magnesium atom is 24.305 amu or (2) the average mass of a mole of magnesium atoms is 24.305 g;

A similar conclusion follows for all of the other elements.

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C101 Class Notes
Prof.N. De Leon

 Chemical Elements Chemistry Volume 1 and 2 Mole of Molecules

Consider one atom of oxygen and one atom of sulfur, and compare their atomic weights.

Oxygen's atomic weight = 15.999 amu

Sulfur's atomic weight = 32.06 amu

The sulfur atom weighs approximately twice as much as the oxygen atom. (32.06 / 15.99 x 2)

Because the sulfur atom weighs twice as much as an oxygen atom, a one gram sample of oxygen contains twice as many atoms as a one gram sample of sulfur. Thus, a two gram sample of sulfur contains the same number of atoms as a one gram sample of oxygen.

From this previous example, one might suggest that a relationship exists between the weight of a sample and the number of atoms in the sample. In fact, scientists have determined that there is a definite relationship between the number of atoms in a sample and the sample's weight. Experimentation has shown that, for any element, a sample containing the atomic weight in grams contains 6.022 x 1023 atoms. Thus 15.999 grams of oxygen contains 6.022 x 1023 atoms, and 32.06 grams of sulfur contains 6.022 x 1023 atoms. This number (6.022 x 1023)is known as Avogadro's number. The importance of Avogadro's number to chemistry should be clear. It represents the number of atoms in X grams of ~Lny element, where X is the atomic weight of the element. It permits chemists to predict and use exact amounts of elements needed to cause desired chemical reactions to occur.

The Mole

A single atom or a few atoms are rarely encountered. Instead, larger, macroscopic quantities are used to quantify or measure collections of atoms or molecules, such as a glass of water, a gallon of alcohol, or two aspirin. Chemists have introduced a large unit of matter, the mole, to deal with macroscopic samples of matter.

One mole represents a definite number of objects, substances, or particles. (For example, a mole of atoms, a mole of ions, a mole of molecules, and even, theoretically, a mole of elephants.) A mole is defined as the quantity of a pure substance that contains 6.022 x 1023 units (atoms, ions, molecules, or elephants) of that substance. In other words, a mole is Avogadro's number of anything.

For any element, the mass of a mole of that element's atoms is the atomic mass expressed in units of grams. For example, to calculate the mass of a mole of copper atoms, simply express the atomic mass of copper in units of grams. Because the atomic mass of copper is 63.546 amu, a mole of copper has a mass of 63.546 grams. The value for the atomic mass of gold is 196.967 amu. Therefore, a mole of gold has a mass of 196.967 grams. The mass of a mole of atoms is called the gram atomic weight (GAW). The mole concept allows the conversion of grams of a substance to moles and vice versa.

Figure 2 A Mole of Gold Compared to a Mole of Copper

Figure 2 contains a ball of gold and a ball of copper.The two balls are of different masses and different sizes, but each contains an identical number of atoms.

Example 1:

A silver bar has a mass of 1870 grams. How many moles of silver are in the bar?

Solution:

Since the atomic mass of silver (Ag) is 107.87 amu, one mole of silver has a mass 107.87 grams. Therefore, there is one mole of Ag per 107.87 grams of Ag

. There are 1870 grams of silver.

Example 2:

Mercury (Hg) is the only metal that exists as a liquid at room temperature. It iszsed in thermometers. A thermometer contains 0.004 moles of mercury. How many grams 6 mercury are in the thermometer?

Solution:

Since the atomic mass of Hg is 201 amu, one mole of Hg has a mass of 201 grams of Hg or . There are 0.004 moles of Hg. 1 mole Hg