CHEMICAL BONDS IN METALS AND NON-METALS

There is no denying the fact that Forces, Electrons, and Bonds in CHEMISTRY CATEGORIES where atoms are the basic building blocks of all types of matter. Atoms link simultaneously to other atoms through chemicals bonds resulting from the strong attractive forces that exist between the atoms.


It is a region that forms when electrons from different atoms interact with each other. The electrons that participate in chemical bonds are the valence electrons, which are the electrons found in an atom’s outermost shell. When two atoms approach each other these outer electrons interact. Electrons repel each other, yet they are attracted to the protons within atoms. The interplay of forces results in some atoms forming bonds with each other and sticking together.
The two main types of bonds formed between atoms are ionic bonds and covalent bonds. An ionic bond is formed when one atom accepts or donates one or more of its valence electrons to another atom. A covalent bond is formed when atoms share valence electrons. The atoms do not always share the electrons equally, so a polar covalent bond (polar covalent bond is such where the atoms do not share the electrons equally) may be the result. When electrons are shared by two metallic atoms a metallic bond may be formed. In a covalent bond, electrons are shared between two atoms. The electrons that participate in metallic bonds may be shared between any of the metal atoms in the region.
• If the electro negativity values of two atoms are similar.
• Metallic bonds form between two metal atoms.
• Covalent bonds form between two non-metal atoms.
• Non polar covalent bonds form when the electro negativity values are very similar.
Carbon Compounds
There are more carbon compounds than there are compounds of all other elements combined. The study of carbon compounds, both natural and synthetic, is called organic chemistry. Plastics, foods, textiles, and many other common substances contain carbon. Hydrocarbon fuels (e.g., natural gas), marsh gas, and the gases resulting from the combustion of fuels (e.g., carbon monoxide and carbon dioxide) are compounds of carbon. With oxygen and a metallic element, carbon forms many important carbonates, such as calcium carbonate (limestone) and sodium carbonate (soda). Certain active metals react with it to make industrially important carbides, such as silicon carbide (an abrasive known as carborundum), calcium carbide, used for producing acetylene gas, and tungsten carbide, an extremely hard substance used for rock drills and metalworking tools.
Polar covalent bonds form when the electro negativity values are a little further apart different.


The Mole and Avogadro’s Constant


Table of Contents
1. Introduction
2. Applications of the Mole
3. Practice Problems
4. Answers to Practice Problems


THE MOLE
The mole, abbreviated mol, is an SI unit which measures the number of particles in a specific substance. One mole is equal to 6.02214179 × 1023 atoms, or other elementary units such as molecules.
Introduction
The number of moles in a system can be determined using the atomic mass of an element, which can be found on the periodic table. This mass is usually an average of the abundant forms of that element found on earth. An element’s mass is listed as the average of all its isotopes on earth.


Example 1
One mole of oxygen atoms contains 6.02214179×1023 oxygen atoms.
The number 6.02214179×1023 alone is called Avogadro’s number or Avogadro’s constant, after the 19th century scientist Amedeo Avogadro.
Each carbon-12 atom weighs about 1.99265 × 1023 g
therefore, (1.99265×1023 g × 6.02214179×1023 atoms) = 12 g of carbon-12.
Applications of the Mole
The mass of a mole of substance is called the molar mass of that substance. The molar mass is used to convert grams of a substance to moles and is used often in chemistry. The molar mass of an element is found on the periodic table, and it is the element’s atomic weight in grams/mole (g/mol).
If the mass of a substance is known, the number of moles in the substance can be calculated. Converting the mass, in grams, of a substance to moles requires a conversion factor of (one mole of substance/molar mass of substance).
The mole concept is also applicable to the composition of chemical compounds. For instance, consider methane, CH4. This molecule and its molecular formula indicate that per mole of methane there is 1 mole of carbon and 4 moles of hydrogen. In this case, the mole is used as a common unit that can be applied to a ratio as shown below:
2 mol H + 1 mol O = 1 mol H2O
The moles of H and O describe the number of atoms of each element that react to form 1 mol of H2O.
To think about what a mole means, one should relate it to quantities such as dozen or pair. Just as a pair can mean two shoes, two books, two pencils, two people, or two of anything else, a mole means 6.02214179×1023 of anything.
Using the following relation:
1 mole = 6.02214179×1023 is analogous to saying:
1 Dozen = (12 eggs)
It is quite difficult to visualize a mole of something because Avogadro’s constant is extremely large. For instance, consider the size of one single grain of wheat. If all the people who have existed in Earth’s history did nothing but count individual wheat grains for their entire lives, the total number of wheat grains counted would still be much less than Avogadro’s constant; the number of wheat grains produced throughout history does not even approach Avogadro’s Number.


How many moles of potassium (K) atoms are in 3.04 grams of pure potassium metal?
SOLUTION
In this example, multiply the mass of K by the conversion factor:
1 mol K / 39.10 grams K
39.10 grams is the molar mass of one mole of K;
cancel out grams, leaving the moles of K:
3.04 g K × (1 mol K/ 39.10 g K) = 0.0778 mol K
Similarly, if the moles of a substance are known, the number grams in the
substance can be determined. Converting moles of a substance to grams requires
a conversion factor of molar mass of substance/one mole of substance.
One simply needs to follow the same method but in the opposite direction.


Example 2
How many grams are 10.78 moles of Calcium (Ca)?
SOLUTION
10.78 mol Ca × (40.08 g Ca/ 1 mol Ca) = 432.1 g Ca
Multiply moles of Ca by the conversion factor 40.08 g Ca/ 1 mol Ca, with
40.08 g being the molar mass of one mole of Ca, which then allows the
cancelation of moles, leaving grams of Ca.
The total number of atoms in a substance can also be determined by
using the relationship between grams, moles, and atoms. If given the mass
of a substance and asked to find the number of atoms in the substance, one
must first convert the mass of the substance, in grams, to moles, as in
Example 1. Then the number of moles of the substance must be converted to atoms.
Converting moles of a substance to atoms requires a conversion factor of
Avogadro’s constant (6.02214179×1023) / one mole of substance.
Verifying that the units cancel properly is a good way to make sure the correct method is used.
Example 3
How many atoms are in a 3.5 g sample of sodium (Na)?
SOLUTION
3.5 g Na × (1 mol Na/ 22.98 g Na) = .152 mol Na
0.152 mol Na × (6.02214179×1023 atoms Na/ 1 mol Na) = 9.15×1022 atoms of Na
In this example, multiply the grams of Na by the conversion
factor 1 mol Na/ 22.98 g Na, with 22.98g being the molar mass of
one mole of Na, which then allows cancelation of grams, leaving moles
of Na. Then, multiply the number of moles of Na by the conversion factor 6.02214179×1023 atoms Na/ 1 mol Na, with 6.02214179×1023 atoms
being the number of atoms in one mole of Na (Avogadro’s constant),
which then allows the cancelation of moles, leaving the number of atoms of Na.


Applications
Using Avogadro’s constant, it is also easy to calculate the number of atoms or molecules present in a substance. By multiplying the number of moles by Avogadro’s constant, the mol units cancel out, leaving the number of atoms. The following table provides a reference for the ways in which these various quantities can be manipulated:
Known Information Multiply By Result
Mass of substance (g) Molar mass (g/mol) Moles of substance
Moles of substance (mol) Avogadro’s constant (atoms/mol) Atoms (or molecules)
Mass of substance (g) Molar mass ((mol/g) × Avogadro’s constant (atoms/mol)) Atoms (or molecules)


Example 1A
Convert to moles: 3.00 grams of potassium (K)
SOLUTION
= 3.00 g K × (1 mol K/ 39.10 g K) = 0.076726 mol K
In this example, multiply the mass of K by the conversion factor:
1 mol K / 39.10 grams K
39.10 grams is the molar mass of one mole of K. Grams can be canceled,
leaving the moles of K.
Example 1B
Convert to grams: 10.00 moles of calcium (Ca)
SOLUTION
This is the calculation in Example 1A performed in reverse.
Multiply moles of Ca by the conversion factor 40.08 g Ca/ 1 mol Ca,
With 40.08 g being the molar mass of one mole of Ca. The moles cancel,
leaving grams of Ca:
10.00 mol Ca × (40.08 g Ca/ 1 mol Ca) = 400.8 grams of Ca
The number of atoms can also be calculated using
Avogadro’s Constant (6.02214179×1023) / one mole of substance.


Example 1C
How many atoms are in a 3.0 g sample of sodium (Na)?
SOLUTION
3.0 g Na × (1 mol Na/ 22.98 g Na) = .130548 mol Na
0.130548 mol Na × (6.02214179×1023 atoms Na/ 1 mol Na) = 7.86179×1022 atoms of Na
Practice Problems
1) Using a periodic table, give the molar mass of the following:
a. H
b. Se
c. Ne
d. Cs
e. Fe


For problems 2-4, convert to moles and find the total number of atoms.
2) 5.06 grams of oxygen 3) 2.14 grams of K 4) 0.134 kg of Li
Convert the following to grams
5) 4.5 moles of C 6) 7.1 moles of Al 7) 2.2 moles of Mg
How many moles are in the product of the reaction
8) 6 mol H + 3 mol O →? mol H2O
9) 1 mol Cl + 1 mol Cl →? mol Cl2
10) 5 mol Na + 4 mol Cl → ? mol NaCl


————————————————————————————————–
Answers to Practice Problems
1a. 1.008 g/mol
1b. 78.96 g/mol
1c. 20.18 g/mol
1d. 132.91g/mol
1e. 55.85 g/mol
2. 5.06g O (1mol/16.00g)= 0.316 mol of O
0.316 mols (6.022×1023 atoms/ 1mol) = 1.904×1023 atoms of O
3. 2.14g K (1mol/39.10g)= 0.055 mol of K
0.055 mols (6.022×1023 atoms/ 1mol) = 3.312×1022 atoms of K
4. 0.134kg Li (1000g/1kg)= 134g Li (1mol/6.941g)= 19.3 mols Li
19.3 (6.022×1023 atoms/ 1mol) = 1.16×1025 atoms of Li
5. 4.5 mols of C (12.011g/1mol) = 54.05 g of C
6. 7.1 mols of Al (26.98g/1mol) = 191.56 g of Al
7. 2.2 mols of Mg (24.31g/1mol) = 53.48 g of MG
8. 6 mol H + 3 mol O → 3 mol H20
9. 1 mol Cl + 1 mol Cl → 1 mol Cl2
10. 5 mol Na + 4 mol Cl → 4 mol NaCl + 1 mol Na (excess)
EMPIRICAL AND MOLECULAR FORMULA


This article is about analytical chemistry. For observation rather than theory, see Empirical relationship.
In chemistry, the empirical formula of a chemical compound is the simplest positive integer ratio of atoms present in a compound.[1] A simple example of this concept is that the empirical formula of hydrogen peroxide, or H2O2, would simply be HO.
An empirical formula makes no mention towards arrangement or number of atoms. It is standard for a lot of ionic compounds, like CaCl2, and for macromolecules, such as SiO2.
The molecular formula on the other hand shows the number of each type of atom in a molecule, also the structural formula shows the arrangement of the molecule. It is possible for different types of compounds to have equal empirical formulas.
Examples
• Glucose (C6H12O6), ribose (C5H10O5), acetic acid (C2H4O2), and formaldehyde (CH2O) all have different molecular formulas but the same empirical formula: CH2O. This is the actual molecular formula for formaldehyde, but acetic acid has double the number of atoms, ribose has five times the number of atoms, and glucose has six times the number of atoms.
• The chemical compound n-hexane has the structural formula CH3CH2CH2CH2CH2CH3, which shows that it has 6 carbon atoms arranged in a chain, and 14 hydrogen atoms. Hexane’s molecular formula is C6H14, and its empirical formula is C3H7, showing a C:H ratio of 3:7.
Calculation
Suppose you are given a compound such as methyl acetate, a solvent commonly used in paints, inks, and adhesives. When methyl acetate was chemically analyzed, it was discovered to have 48.64% carbon (C), 8.16% hydrogen (H), and 43.20% oxygen (O). For the purposes of determining empirical formulas, we assume that we have 100 g of the compound. If this is the case, the percentages will be equal to the mass of each element in grams.
Step 1
Change each percentage to an expression of the mass of each element in grams. That is, 48.64% C becomes 48.64 g C, 8.16% H becomes 8.16 g H, and 43.20% O becomes 43.20 g O.
Step 2
Convert the amount of each element in grams to its amount in moles.


Step 3
Divide each of the found values by the smallest of these values (2.7)


Step 4
If necessary, multiply these numbers by integers in order to get whole numbers; if an operation is done to one of the numbers, it must be done to all of them.


Thus, the empirical formula of methyl acetate is C3H6O2. This formula also happens to be methyl acetate’s molecular formula.


Problem


The simplest formula for vitamin C is C3H4O3. Experimental data indicates that the molecular mass of vitamin
C is about 180. What is the molecular formula of vitamin C?


Solution


First, calculate the sum of the atomic masses for C3H4O3. Look up the atomic masses for the elements from the Periodic Table. The atomic masses are found to be: H is 1.01, C is 12.01, O is 16.00
Plugging in these numbers, the sum of the atomic masses for C3H4O3 is:
3(12.0) + 4(1.0) + 3(16.0) = 88.0
This means the formula mass of vitamin C is 88.0. Compare the formula mass (88.0) to the approximate molecular mass (180). The molecular mass is twice the formula mass (180/88 = 2.0), so the simplest formula must be multiplied by 2 to get the molecular formula:
molecular formula vitamin C = 2 x C3H4O3 = C6H8O6.
answer:C6H8O6


An approximate molecular mass is usually sufficient to determine the formula mass, but the calculations tend not to work out ‘even’ as in this example. You are looking for the closest whole number to multiply by the formula mass to get the molecular mass.
In chemistry and physics, the Avogadro constant (symbols: L, NA) is defined as the number of constituent particles (usually atoms or molecules) per mole of a given substance, where the mole (abbreviation: mol) is one of the seven base units in the International System of Units (SI). The Avogadro constant has dimensions of reciprocal mol and its value is equal to 6.02214129(27)×1023 mol−1. The constant happens to be quite close to an integer power of two, specifically only about 0.37% less than 279 mol−1, making the latter a useful approximation in nuclear physics when considering chain reaction growth rates.[citation needed]
Previous definitions of chemical quantity involved Avogadro’s number, a historical term closely related to the Avogadro constant but defined differently: Avogadro’s number was initially defined by Jean Baptiste Perrin as the number of atoms in one gram-molecule of atomic hydrogen, meaning (in modern terminology) one gram of (atomic) hydrogen. It was later redefined as the number of atoms in 12 grams of the isotope carbon-12 and still later generalized to relate amounts of a substance to their molecular weight.[4] For instance, to a first approximation, 1 gram of hydrogen, which has a mass number of 1 (atomic number 1), has 6.023×1023 hydrogen atoms. Similarly, 12 grams of carbon 12, with the mass number of 12 (atomic number 6), has the same number of carbon atoms, 6.023×1023. Avogadro’s number is a dimensionless quantity and has the numerical value of the Avogadro constant given in base units.
The Avogadro constant is fundamental to understanding both the makeup of molecules and their interactions and combinations. For instance, since one atom of oxygen will combine with two atoms of hydrogen to create one molecule of water (H2O), one can similarly see that one mole of oxygen (6.022×1023of O atoms) will combine with two moles of hydrogen (2 × 6.022×1023 of H atoms) to make one mole of H2O.
Mole and moles are frequently abbreviated as mol in chemical and mathematic notation.
Revisions in the base set of SI units necessitated redefinitions of the concepts of chemical quantity and so Avogadro’s number, and its definition, was deprecated in favor of the Avogadro constant and its definition. Changes in the SI units are proposed that will precisely fix the value of the constant to exactly 6.02214X×1023 when it is expressed in the unit mol−1 (see New SI definitions, in which an “X” at the end of a number means one or more final digits yet to be agreed upon).
Value of NA[5] in various units
6.02214129(27)×1023 mol−1
2.73159734(12)×1026 (lb-mol)−1
1.707248434(77)×1025 (oz-mol)−1
History
The Avogadro constant is named after the early 19th century Italian scientist Amedeo Avogadro, who in 1811 first proposed that the volume of a gas (at a given pressure and temperature) is proportional to the number of atoms or molecules regardless of the nature of the gas.[6] The French physicist Jean Perrin in 1909 proposed naming the constant in honor of Avogadro.[7]Perrin won the 1926 Nobel Prize in Physics, largely for his work in determining the Avogadro constant by several different methods.[8]
The value of the Avogadro constant was first indicated by Johann Josef Loschmidt who in 1865 estimated the average diameter of the molecules in air by a method that is equivalent to calculating the number of particles in a given volume of gas.[9] This latter value, the number density of particles in an ideal gas, is now called the Loschmidt constant in his honor, and is related to the Avogadro constant, NA, by


where p0 is the pressure, R is the gas constant and T0 is the absolute temperature. The connection with Loschmidt is the root of the symbol L sometimes used for the Avogadro constant, and German language literature may refer to both constants by the same name, distinguished only by the units of measurement.[10]
Accurate determinations of Avogadro’s number require the measurement of a single quantity on both the atomic and macroscopic scales using the same unit of measurement. This became possible for the first time when American physicist Robert Millikan measured the charge on an electron in 1910. The electric charge per mole of electrons is a constant called the Faraday constant and had been known since 1834 when Michael Faraday published his works on electrolysis. By dividing the charge on a mole of electrons by the charge on a single electron the value of Avogadro’s number is obtained.[11] Since 1910, newer calculations have more accurately determined the values for the Faraday constant and the elementary charge. (See below)
Perrin originally proposed the name Avogadro’s number (N) to refer to the number of molecules in one gram-molecule of oxygen (exactly 32g of oxygen, according to the definitions of the period),[7] and this term is still widely used, especially in introductory works.[12] The change in name to Avogadro constant (NA) came with the introduction of the mole as a base unit in the International System of Units (SI) in 1971,[13] which recognized amount of substance as an independent dimension of measurement.[14] With this recognition, the Avogadro constant was no longer a pure number, but had a unit of measurement, the reciprocal mole (mol−1).[14]
While it is rare to use units of amount of substance other than the mole, the Avogadro constant can also be expressed in units such as the pound mole (lb-mol) and the ounce mole (oz-mol).
NA = 2.73159757(14)×1026 (lb-mol)−1 = 1.707248479(85)×1025 (oz-mol)−1
General role in science
Avogadro’s constant is a scaling factor between macroscopic and microscopic (atomic scale) observations of nature. As such, it provides the relation between other physical constants and properties. For example, it establishes a relationship between the gas constant R and the Boltzmann constant kB,


and the Faraday constant F and the elementary charge e,


The Avogadro constant also enters into the definition of the unified atomic mass unit, u,


where Mu is the molar mass constant.
Measurement
Coulometry
The earliest accurate method to measure the value of the Avogadro constant was based on coulometry. The principle is to measure the Faraday constant, F, which is the electric charge carried by one mole of electrons, and to divide by the elementary charge, e, to obtain the Avogadro constant.


The classic experiment is that of Bower and Davis at NIST,[15] and relies on dissolving silver metal away from the anode of an electrolysis cell, while passing a constant electric current I for a known time t. If m is the mass of silver lost from the anode and Ar the atomic weight of silver, then the Faraday constant is given by:


The NIST scientists devised a method to compensate for silver lost from the anode by mechanical causes, and conducted an isotope analysis of the silver used to determine its atomic weight. Their value for the conventional Faraday constant is F90 = 96,485.39(13) C/mol, which corresponds to a value for the Avogadro constant of 6.0221449(78)×1023 mol−1: both values have a relative standard uncertainty of 1.3×10−6.
Electron mass measurement[edit]
The Committee on Data for Science and Technology (CODATA) publishes values for physical constants for international use. It determines the Avogadro constant[16] from the ratio of the molar mass of the electron Ar(e)Mu to the rest mass of the electron me:


The relative atomic mass of the electron, Ar(e), is a directly-measured quantity, and the molar mass constant, Mu, is a defined constant in the SI. The electron rest mass, however, is calculated from other measured constants:[16]


As may be observed in the table of 2006 CODATA values below,[17] the main limiting factor in the precision of the Avogadro constant is the uncertainty in the value of the Planck constant, as all the other constants that contribute to the calculation are known more precisely.
Constant Symbol 2006 CODATA value Relative standard uncertainty Correlation coefficient
with NA
Electron relative atomic mass Ar(e) 5.485 799 0943(23)×10–4 4.2×10–10 0.0082
Molar mass constant
Mu 0.001 kg/mol = 1 g/mol Defined —
Rydberg constant
R∞ 10 973 731.568 527(73) m−1 6.6×10–12 0.0000
Planck constant
h 6.626 068 96(33)×10–34 J s 5.0×10–8 −0.9996
Speed of light
c 299 792 458 m/s Defined —
Fine structure constant
α 7.297 352 5376(50)×10–3 6.8×10–10 0.0269
Avogadro constant NA 6.022 141 79(30)×1023 mol−1 5.0×10–8 1
X-ray crystal density (XRCD) methods[edit]


Ball-and-stick model of the unit cell of silicon. X-ray diffraction measures the cell parameter, a, which is used to calculate a value for Avogadro’s constant.
A modern method to determine the Avogadro constant is the use of X-ray crystallography. Silicon single crystals may be produced today in commercial facilities with extremely high purity and with few lattice defects. This method defines the Avogadro constant as the ratio of the molar volume, Vm, to the atomic volume Vatom:
where and n is the number of atoms per unit cell of volume Vcell.
The unit cell of silicon has a cubic packing arrangement of 8 atoms, and the unit cell volume may be measured by determining a single unit cell parameter, the length of one of the sides of the cube, a.[18]
In practice, measurements are carried out on a distance known as d220(Si), which is the distance between the planes denoted by the Miller indices {220}, and is equal to a/√8. The 2006 CODATA value for d220(Si) is 192.0155762(50) pm, a relative uncertainty of 2.8×10−8, corresponding to a unit cell volume of 1.60193304(13)×10−28 m3.
The isotope proportional composition of the sample used must be measured and taken into account. Silicon occurs in three stable isotopes (28Si, 29Si,30Si), and the natural variation in their proportions is greater than other uncertainties in the measurements. The atomic weight Ar for the sample crystal can be calculated, as the relative atomic masses of the three nuclides are known with great accuracy. This, together with the measured density ρ of the sample, allows the molar volume Vm to be determined:


where Mu is the molar mass constant. The 2006 CODATA value for the molar volume of silicon is 12.058 8349(11) cm3mol−1, with a relative standard uncertainty of 9.1×10−8.[19]
As of the 2006 CODATA recommended values, the relative uncertainty in determinations of the Avogadro constant by the X-ray crystal density method is 1.2×10−7, about two and a half times higher than that of the electron mass method.
One of the master opticians at the Australian Centre for Precision Optics (ACPO) holding a one-kilogram single-crystal silicon sphere for the International Avogadro Coordination.
The International Avogadro Coordination (IAC), often simply called the “Avogadro project”, is a collaboration begun in the early 1990s between various national metrology institutes to measure the Avogadro constant by the X-ray crystal density method to a relative uncertainty of 2×10−8 or less.[20] The project is part of the efforts to redefine the kilogram in terms of a universal physical constant, rather than the International Prototype Kilogram, and complements the measurements of the Planck constant using watt balances.[21][22] Under the current definitions of the International System of Units (SI), a measurement of the Avogadro constant is an indirect measurement of the Planck constant:


The measurements use highly polished spheres of silicon with a mass of one kilogram. Spheres are used to simplify the measurement of the size (and hence the density) and to minimize the effect of the oxide coating that inevitably forms on the surface. The first measurements used spheres of silicon with natural isotopic composition, and had a relative uncertainty of 3.1×10−7.[23][24][25] These first results were also inconsistent with values of the Planck constant derived from watt balance measurements, although the source of the discrepancy is now believed to be known.[22]
The main residual uncertainty in the early measurements was in the measurement of the isotopic composition of the silicon to calculate the atomic weight so, in 2007, a 4.8-kg single crystal of isotopically-enriched silicon (99.94% 28Si) was grown,[26][27] and two one-kilogram spheres cut from it. Diameter measurements on the spheres are repeatable to within 0.3 nm, and the uncertainty in the mass is 3 µg. Full results from these determinations were expected in late 2010.[28] Their paper, published in January 2011, summarized the result of the International Avogadro Coordination and presented a measurement of the Avogadro constant to be 6.02214078(18)×1023 mol−1
Avogadro’s Number in Practice
How on Earth did chemists settle on such a seemingly arbitrary figure for Avogadro’s number? To understand how it was derived, we have to first tackle the concept of the atomic mass unit (amu). Theatomic mass unit is defined as 1/12 of the mass of one atom of carbon-12 (the most common isotope of carbon). Here’s why that’s neat: Carbon-12 has six protons, six electrons and six neutrons, and because electrons have very little mass, 1/12 of the mass of one carbon-12 atom is very close to the mass of a single proton or a single neutron. The atomic weights of elements (those numbers you see below the elements on the periodic table) are expressed in terms of atomic mass units as well. For instance, hydrogen has, on average, an atomic weight of 1.00794 amu.
Unfortunately, chemists don’t have a scale that can measure atomic mass units, and they certainly don’t have the ability to measure a single atom or molecule at a time to carry out a reaction. Since different atoms weigh different amounts, chemists had to find a way to bridge the gap between the invisible world of atoms and molecules and the practical world of chemistry laboratories filled with scales that measure in grams. In order to do this, they created a relationship between the atomic mass unit and the gram, and that relationship looks like this:
1 amu = 1/6.0221415 x 1023 grams
This relationship means that if we had Avogadro’s number, or one mole, of carbon-12 atoms (which has an atomic weight of 12 amu by definition), that sample of carbon-12 would weigh exactly 12 grams. Chemists use this relationship to easily convert between the measurable unit of a gram and the invisible unit of moles, of atoms or molecules.
Now that we know how Avogadro’s number comes in handy, we need to examine one last question: How did chemists determine how many atoms are in a mole in the first place? The first rough estimate came courtesy of physicist Robert Millikan, who measured the charge of an electron. The charge of a mole of electrons, called a Faraday, was already known by the time Millikan made his discovery.
Dividing a Faraday by the charge of an electron, then, gives us Avogadro’s number. Over time, scientists have found new and more accurate ways of estimating Avogadro’s number, most recently using advanced techniques like using X-rays to examine the geometry of a 1 kilogram sphere of silicon and extrapolating the number of atoms it contained from that data. And while the kilogram is the basis for all units of mass, some scientists want to begin using Avogadro’s number instead, much the way we now define the length of a meter based on the speed of light instead of the other way around.
Atoms and
Molecules


The basic building blocks of the “normal” matter that we see in the Universe are atoms, and combinations of atoms that we call molecules. We first consider atoms and then molecules. However, we shall see that although “normal matter” is composed of atoms and molecules, most of the matter in the Universe is not in the form of atoms or molecules, but rather in the form of a plasma. We discuss plasmas in the next section.
Constituents of Atoms
Atoms are composed of three classes of constituents, as illustrated in the following table.
Constituent Symbol Charge Mass
Electrons e- -1 9.1 x 10-28 g
Protons p+ +1 1836 x electron mass
Neutrons n 0 Approximately that of p+


Thus, most of the mass of atoms resides in the neutrons and protons which occupy the dense central region called the nucleus (see the Bohr atom below).
The number of protons (or the number of electrons) is called the atomic number for the atom. The total number of protons plus neutrons is called the atomic mass number for the atom. Atoms are electrically neutral because the number of negatively-charged electrons is exactly equal to the number of positively-charged protons. The number of neutrons is approximately equal to the number of protons for stable light nuclei, and is about 1-2 times the number of protons for the heavier stable nuclei.
Isotopes of an Element
Atoms having the same number of protons (and therefore the same number of electrons) but different numbers of neutrons are called isotopes of the element in question. Thus, the isotopes of an element have the same atomic number but differ in their atomic mass number. A compact notation for isotopes of an element is illustrated by the following examples.


In this notation the element is represented by its chemical symbol, the atomic number is denoted by a lower left subscript, the number of neutrons is denoted by a lower right subscript, and the atomic mass number is denoted by an upper left superscript (some of these superscripts and subscripts may be omitted, depending on the context).
Thus, the above symbols denote, respectively, the mass-235 and mass-238 isotopes of uranium (symbol U), and the mass-1,-2,and -3 isotopes of hydrogen (symbol H). The mass-2 isotope of hydrogen is also called deuterium and the mass-3 isotope is also called tritium.
Periodic Table of the Elements
The elements have properties that repeat themselves periodically with variation of the number of electrons (atomic number). A chart of the elements arranged to show this periodicity is termed a periodic table (of the elements). Here is a periodic table of the elements which gives the atomic number and symbol for all elements, and the name and basic chemical properties for each of these elements if you click on the element’s symbol in the resulting table.