Chemistry
Relative Atomic Mass
Relative atomic mass is the average mass of an element's atoms compared to the mass of a carbon-12 atom. It is expressed in atomic mass units (u) and is calculated by taking into account the abundance of each isotope of the element. This value is used to compare the masses of different atoms in the periodic table.
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7 Key excerpts on "Relative Atomic Mass"
eBook - PDFFoundations of Chemistry
An Introductory Course for Science Students
- Philippa B. Cranwell, Elizabeth M. Page(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
In making these new substances, scientists need to know how much of one molecule or com-pound will react with another and how much product will be obtained as a result. To achieve this, we need to be able to count molecules or groups of mole-cules. And to use the right amount of reacting substance, we need to be able to measure it by determining its mass or volume or by knowing its concentration. This means we need to know about the relative masses of the atoms and mole-cules in reacting substances. 3.1.1 Relative Atomic Mass, A r In Chapter 1, we saw that different atoms have different numbers of neutrons, protons, and electrons, and this means that atoms of different elements have dif-ferent masses. Because atoms are so small, it is impossible to weigh just one, so we measure their masses relative to each other. The stable isotope of carbon-12, 12 6 C, has been chosen as the standard by which to compare the masses of other atoms. One atom of carbon-12 has been assigned a mass of exactly 12 atomic mass units (amu). So one amu is equal to one-twelfth the mass of an atom of carbon-12, which is equivalent to 1.66 × 10 -27 kg. The masses of other atoms are obtained by comparing the average mass of one atom of the element to one-twelfth the mass of a carbon-12 atom. This value is called the Relative Atomic Mass of the element and has the symbol, A r . A r = average mass of one atom of element 1 12 mass of one atom of carbon-12 The Relative Atomic Mass of an element, A r , is defined as the average mass of one atom of the element compared to the mass of one-twelfth of an atom of carbon-12. Using this scale, the Relative Atomic Masses as shown in a periodic table of the elements are obtained. Hydrogen, the lightest element, has a Relative Atomic Mass of 1.008 when given to three decimal places (four significant figures). Nitrogen, which is slightly heavier than carbon, has a A r of 14.007.
eBook - ePub- Nigel P. Freestone(Author)
- 2016(Publication Date)
- Bentham Science Publishers(Publisher)
Relative Atomic & Formula Mass, Percentage Composition, Empirical & Molecular Formulae Nigel P. FreestoneKeywords: Atoms, Average Relative Atomic Mass, Calculations, Empirical and molecular formula, Isotopes, Percentage composition, Relative Atomic Mass (Ar ), Relative formula mass (Mr ).1.1. Relative Atomic Mass (AR )Chemical elements are the building blocks from which everything is constructed, from specks of dust to mobile phones and from flora and fauna to the clothes we wear. There are over a 100 known elements. An element is a pure substance that cannot be chemically broken down. The smallest unit of an element is the atom. Different atoms have different masses. The mass of an atom is so small that it is more convenient to compare atom masses, rather than refer to their actual mass. The standard for this relative scale is an atom of carbon-12, which has a Relative Atomic Mass (Ar ) of 12.The Table below lists the Relative Atomic Mass (Ar ) values of some common elements.Selected Relative Atomic Mass Values
Element Approximate Relative Atomic Mass (Ar ) Hydrogen 1 Carbon 12 Nitrogen 14 Oxygen 16 Sodium 23 Magnesium 24 Silicon 28 Calcium 40 Bromine 80 These Relative Atomic Mass values tell us for example that sodium atoms (Ar = 23) are 23 times heavier than hydrogen atoms (Ar = 1), two atoms of neon (Ar = 20) have the same mass as one atom of calcium (Ar = 40) and that three oxygen atoms (Ar = 16 ) weigh the same as two magnesium atoms (Ar = 24).Relative Atomic Masses are listed in the periodic table.Isotopes
Elements are defined by their proton (atomic) number. An atom with 7 protons is always nitrogen (N), an atom with 20 protons is always calcium and an atom with 70 protons must therefore always be gold (Au). Isotopes are atoms that have the same number of protons, but have different numbers of neutrons. For example, carbon has three naturally occurring isotopes, often referred to as simply carbon-12, carbon-13 and carbon-14, with Relative Atomic Masses of 12, 13 and 14, respectively. Since carbon has a proton number of 6, the isotopes contain 6, 7 and 8 neutrons, respectively.
eBook - PDF- Joyce James, Colin Baker, Helen Swain(Authors)
- 2008(Publication Date)
- Wiley-Blackwell(Publisher)
Relative Atomic Mass An atom of the carbon-12 isotope was given a Relative Atomic Mass of 12.0 000 u (u = atomic Mass number (number of protons plus neutrons) 23 Atomic number 11 Na Sodium has 11 protons and 11 electrons and 12 neutrons. (23 − 11 = 12) Fig. 1.8 Atomic and mass numbers. Activity Mercury is a toxic element and can lead to death. In hospitals it is commonly used in thermometers and sphygmomanometers. Look up the symptoms of mercury poisoning and the type of treatment given to such a patient. mass unit). The relative masses of all other atoms were then obtained by comparison to this standard atom. The Relative Atomic Mass is also called ‘atomic mass’, ‘average atomic mass’ and in the past ‘atomic weight’. Most elements contain a mixture of isotopes. The Relative Atomic Mass of an element is the average mass of one atom, taking account of all its isotopes and their relative proportions compared with an atom of carbon-12. For example chlorine consists of two isotopes with mass numbers of 35 and 37. These iso-topes are written as chlorine-35 and chlorine-37 respectively. Naturally occurring chlorine contains 75% of chlorine-35 and 25% of chlo-rine-37. The Relative Atomic Mass of chlorine is: Most Relative Atomic Masses are not whole numbers because most elements have many isotopes (Table 1.3). Electron arrangement Electrons are arranged in shells around the nucleus. The number of electrons in a neutral atom is given by the atomic number. The rules for filling the shells are: 75 35 25 37 100 35 5 ¥ ( ) + ¥ ( ) = . u 1. The shell nearest the nucleus fills first 2. Only a given number of electrons are allowed in each shell Examples 1 st shell 2 2 nd shell 8 3 rd shell 18 The diagrams like those in Figs 1.9 and 1.10 are quite cumbersome to draw each time, so a shorthand form is often used to show how elec-trons are arranged – this is called the electronic configuration.
eBook - PDFChemistry for Today
General, Organic, and Biochemistry
- Spencer Seager, Michael Slabaugh, Maren Hansen, , Spencer Seager, Spencer Seager, Michael Slabaugh, Maren Hansen(Authors)
- 2021(Publication Date)
- Cengage Learning EMEA(Publisher)
Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. Atoms and Molecules 53 Thus, we see that the relative masses used in the periodic table give the same results as the actual masses, when the masses of different atoms are compared with one another. Actual atomic masses are given in mass units such as grams, but the relative values used in the periodic table are given in units referred to as atomic mass units. The symbol “u” will be used to represent atomic mass units throughout the text. The actual mass represented by a single atomic mass unit is 1 12 of the mass of a single carbon-12 atom or 1.661 3 10 224 g. The relative masses of the elements as given in the periodic table are referred to as atomic masses or atomic weights. We will use the term atomic weight in this book. In those cases where the naturally occurring element exists in the form of a mixture of iso- topes, the recorded atomic weight is the average value for the naturally occurring isotope mixture. This idea is discussed in Section 2.5. A convenient way to visualize the concept of relative masses, and to compare the relative masses of atoms, involves the use of the simple child’s toy called a seesaw or teeter-totter, shown below. When the masses on both sides of the central pivot are equal, the seesaw will be in balance. When the masses on each side are different, the seesaw will be out of balance, and the side atomic mass unit (u) A unit used to express the relative masses of atoms. One u is equal to 1 12 the mass of an atom of carbon-12. atomic weight The mass of an average atom of an element expressed in atomic mass units.
- Barron's Educational Series, Albert S. Tarendash(Authors)
- 2021(Publication Date)
- Barrons Educational Services(Publisher)
Pages 83–84 ) 4.2Concept: The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes. (Pages 83–84 )Skills:- Given the atomic mass of an element, determine the mass number of the element’s most abundant isotope. (Pages 83–84 )
- Calculate the atomic mass of an element, given the masses and ratios of the element’s naturally occurring isotopes. (Pages 83–85 )
Concept: The empirical formula of a compound is the simplest whole number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound. (Pages 88–89 )Skills:- Determine the molecular formula of a compound, given the empirical formula and the molecular mass. (Pages 88–89 )
- Determine the empirical formula from the molecular formula. (Pages 88–89 )
Concept: The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram formula mass) of a substance is the mass of 1 mole of that substance. (Pages 85–86 )Skills:- Calculate the formula mass and the molar mass (gram-formula mass) of a substance. (Pages 87–88 )
- Determine the number of moles of a substance, given its mass. (Pages 85–86 )
- Determine the mass of a given number of moles of a substance. (Pages 85–86 )
Concept: The percent composition by mass of each element in a compound can be calculated mathematically. (Pages 87–88 )Skill: Determine the percent composition(s) of one or more elements in a compound. (Pages 87–88 )4.6Concept: A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. (Pages 90–91 )Skill: Solve simple mole-mole stoichiometry problems, given a balanced equation. (Pages 90–91
eBook - PDF- Bernard Moody(Author)
- 2013(Publication Date)
- Arnold(Publisher)
Calculated Relative Atomic Mass (11 x 1.008) + (12 x 1.009) = 23.196 Relative Atomic Mass by experiment = 22.990 In the symbols 2*He and nNa, the upper figure refers to the nucléon (mass) number (A) and the lower figure to the proton (atomic) number (Z) defined earlier. In Chapter 1, 16 0, 17 0, 18 0 and 12 C were mentioned: the upper figure is sometimes placed after the symbol. The determination of atomic masses by the mass spectrometer Isotopic masses and the abundance ratio of the isotopes of elements may be determined, and hence the weighted mean unitless Relative Atomic Mass calculated (or molar mass/gmol 1 ), taking into account the proportions of the atoms of different mass present. A mass spectograph generally used for individual mass determinations, delivers a photographic record while a mass spectrometer incorporates a meter to measure ion current and is used to determine isotopic abundance ratios. Today mass spectrometry is concerned with the deter-mination of relative molecular mass (molar mass/ g mol 1 ) and the deduction of molecular structures, particularly of organic compounds. The principle of the mass spectrometer is described here. Positive ions are commonly produced by elec-tron bombardment, M + _?e --M + + 2_?e but ultraviolet irradiation, sparking or field emis-sion from a surface upon which the sample has been absorbed may be used. The element, or a suit-able compound, is introduced as a gas under very low pressure (1 x 10 4 Pa (N m 2 )) to avoid inter-ference from other entities as far as possible. Molecules may ionize to ions with one or more charges, or simply disintegrate. The positive ions are accelerated by a powerful electric field, a narrow beam is selected and this enters the field of an electromagnetic analyser. The ions sweep round a semi-circular curved path, the radius of curvature of which is related to the magnitude of the accelerating voltage (ocyfÉ) and to the strength of the magnetic field (
eBook - ePub- Atef Korchef(Author)
- 2022(Publication Date)
- CRC Press(Publisher)
One mole (1 mol) contains 6.022 × 1023 entities (atoms, molecules, ions).The atomic mass (the mass of one atom) is the average mass of all the isotopes of an atom with respect to their abundances. The atomic masses of all elements are given in the periodic table of elements. The atomic mass is commonly expressed in atomic mass units (amu).1 a m u =g = 1.660 ×1 /N Ag10− 24- Example: A calcium (Ca) atom has a mass of 40 amu. The mass of calcium atom in g is (40/NA ) g = 6.642 × 10−23 g
The molecular mass is the sum ofthe atomic masses of all the atoms in the molecule. It is commonly expressed in atomic mass units (amu).- Example: One molecule of sulfuric acid (H2 SO4 ) contains two atoms of H, one atom of S and four atoms of O. The molecular mass M of sulfuric acid is
M = 2 ×+ 1 ×atomic mass of H+ 4 ×atomic mass of Satomic mass of O=+2 × 1+1 × 32= 98 amu4 × 16The mass of one sulfuric acid molecule in g is (98/NA )g = 1.627 × 10−22 gFor monatomic elements, the molar mass (mass of one mole, abbreviated as M or Mwt ) is the numerical value on the periodic table expressed in g mol−1 .- Example: The atomic mass of O is equal to 16 amu and the molar mass of O is equal to 16 gmol−1 .
For molecules, the molar mass is the sum of the molar masses of each of the atoms in the molecular formula.- Example: A molecule of H2 O contains two H atoms and one O atom. The atomic mass of O is 16 amu and the atomic mass of H is 1 amu. The molecular mass of H2 O is equal to 16 + 2 × 1 = 18 amu and the molar mass of H2 O is equal to 18 gmol−1
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