The Structure & Stability of Atoms

Early Atomic History There have been many different theories, reflecting different times and cultures, to explain the composition of matter. There have been many different theories, reflecting different times and cultures, to explain the composition of matter. In addition, chemical reactions, refinements of ores, purification of salt, etc. have been carried out for thousands of years. In addition, chemical reactions, refinements of ores, purification of salt, etc. have been carried out for thousands of years.

Early Atomic History The ancient Greek philosophers theorized that matter is discrete, rather than continuous. The ancient Greek philosophers theorized that matter is discrete, rather than continuous. Some, notably Demokritos, suggested that there is some small unit of matter that still retains the properties of the larger sample. It was thought that these smaller pieces of matter were indivisible, and were given the name atomos from which we get our modern word atoms. Some, notably Demokritos, suggested that there is some small unit of matter that still retains the properties of the larger sample. It was thought that these smaller pieces of matter were indivisible, and were given the name atomos from which we get our modern word atoms.

Early Atomic Theory During the next 2000 years, a lot was learned about matter. Several elements were discovered, metals were refined, acids prepared, etc. During the next 2000 years, a lot was learned about matter. Several elements were discovered, metals were refined, acids prepared, etc. In the mid-1600s, the scientific (rather than the philosophical or applied) study of the nature matter began to take shape. In the mid-1600s, the scientific (rather than the philosophical or applied) study of the nature matter began to take shape.

Early Atomic Theory Since most laboratories contained rudimentary equipment- burners and scales, many experiments involved the measurement of changes in volumes (for gases) and masses during chemical reactions. Since most laboratories contained rudimentary equipment- burners and scales, many experiments involved the measurement of changes in volumes (for gases) and masses during chemical reactions. Based on measurements and observations, several scientific laws were developed. These laws form the basis for our understanding of the composition of matter. Based on measurements and observations, several scientific laws were developed. These laws form the basis for our understanding of the composition of matter.

The Law of Conservation of Mass Antoine Lavoisier ( ) measured the masses of reactants and products for a variety of chemical reactions. He determined that matter is neither created nor destroyed during a chemical reaction. This is known as the law of conservation of mass. Antoine Lavoisier ( ) measured the masses of reactants and products for a variety of chemical reactions. He determined that matter is neither created nor destroyed during a chemical reaction. This is known as the law of conservation of mass.

The Law of Definite Proportion Joseph Proust ( ) determined the chemical composition of many compounds. He found that a given compound always contains the exact same proportion of elements by mass. This is known as the law of definite proportion. Joseph Proust ( ) determined the chemical composition of many compounds. He found that a given compound always contains the exact same proportion of elements by mass. This is known as the law of definite proportion. For example, all samples of water contain 88.8% oxygen by mass, and 11.2% hydrogen by mass.

The Law of Multiple Proportions This chemical law applies when two (or more) elements can combine to form different compounds. This chemical law applies when two (or more) elements can combine to form different compounds. Common examples are carbon monoxide and carbon dioxide, or water and hydrogen peroxide. John Dalton ( ) conducted experiments on these types of compounds, and determined that there is a simple relationship between the masses of one element relative to the others. John Dalton ( ) conducted experiments on these types of compounds, and determined that there is a simple relationship between the masses of one element relative to the others.

The Law of Multiple Proportions When two elements form a series of compounds, the ratios of the masses of one element that combine with a fixed mass of the other element are always in a ratio of small whole numbers. When two elements form a series of compounds, the ratios of the masses of one element that combine with a fixed mass of the other element are always in a ratio of small whole numbers. The meaning of this law is difficult to understand unless it is illustrated using a specific series of compounds.

The Law of Multiple Proportions Consider the compounds of water and hydrogen peroxide. At this point in history, chemists knew the compounds were different, and that they both contain (or can be broken down into) the elements hydrogen and oxygen. They did not yet know the formulas for either compound, nor was the concept of atoms fully developed. Consider the compounds of water and hydrogen peroxide. At this point in history, chemists knew the compounds were different, and that they both contain (or can be broken down into) the elements hydrogen and oxygen. They did not yet know the formulas for either compound, nor was the concept of atoms fully developed.

The Law of Multiple Proportions Analysis of 100 grams of the compounds produced the following data: Analysis of 100 grams of the compounds produced the following data: Compound Mass of oxygen/100g of compound Mass of hydrogen/100g of compound Grams of oxygen/gram of hydrogen water 88.8 grams O 11.2 grams H 7.93 gO/gH Hydrogen peroxide grams O 5.94 grams H 15.8 gO/gH

The Law of Multiple Proportions The Law of Multiple Proportions is illustrated when the numbers in the last column are compared. The Law of Multiple Proportions is illustrated when the numbers in the last column are compared. Compound Grams of oxygen/gram of hydrogen water 7.93 gO/gH Hydrogen peroxide 15.8 gO/gH 15.8/7.93 = 2/1 The small whole number ratio suggests that there is twice as much oxygen in hydrogen peroxide as there is in water.

The Law of Multiple Proportions The key feature is that small whole numbers are generated. The results support the hypothesis that molecules consist of various combinations of atoms, and that atoms are the smallest unit of matter. The ratio doesn’t produce fractions, since there is no such thing as a fraction of an atom. The key feature is that small whole numbers are generated. The results support the hypothesis that molecules consist of various combinations of atoms, and that atoms are the smallest unit of matter. The ratio doesn’t produce fractions, since there is no such thing as a fraction of an atom. For the example cited, we would propose that hydrogen peroxide contains twice as many oxygen atoms/hydrogen atoms than does water. We cannot, however, determine the actual formula of either compound. For the example cited, we would propose that hydrogen peroxide contains twice as many oxygen atoms/hydrogen atoms than does water. We cannot, however, determine the actual formula of either compound.

The Law of Multiple Proportions

Dalton’s Atomic Theory (1808) 1. Each element consists of tiny particles called atoms. 2. The atoms of a given element are identical, and differ from the atoms of other elements. 3. Compounds are formed when atoms of different elements combine chemically. A specific compound always has the same relative number and types of atoms. 4.Chemical reactions involve the reorganization of atoms, or changes in the way they are bound together.

Sub-Atomic Particles The period from approximately involved the study of the nature of the atom, using two relatively new tools: electricity and radioactivity. The period from approximately involved the study of the nature of the atom, using two relatively new tools: electricity and radioactivity. Scientists knew that atoms of different elements had different relative atomic masses and different properties, and they wanted to find out the reasons for the differences. Scientists knew that atoms of different elements had different relative atomic masses and different properties, and they wanted to find out the reasons for the differences.

Sub-Atomic Particles J.J. Thomson ( ) studied the properties of cathode rays. The rays are produced in partially evacuated tubes containing electrodes at either end. J.J. Thomson ( ) studied the properties of cathode rays. The rays are produced in partially evacuated tubes containing electrodes at either end. The rays are invisible, unless a phosphorescent screen is used. The rays are invisible, unless a phosphorescent screen is used.

Sub-Atomic Particles

(Cathode) (Anode) Cathode Rays

Sub-Atomic Particles Thomson made the following observations: 1. The cathode rays had the same properties regardless of the metal used for the cathode. 2. The rays traveled from the cathode (- charged) to the anode (+ charged). 3. The rays were attracted to the positive plate of an external electrical field, and repelled by the negative plate.

Sub-Atomic Particles Thomson concluded: 1. The cathode rays are a stream of negatively charged particles called electrons. 2. All atoms contain electrons, and the electrons from all elements are identical. 3. The atom must also contain matter with a positive charge, as atoms are neutral in charge.

Sub-Atomic Particles Thomson also carried out deflection measurements, in which he applied a magnetic field to deflect the beam along with an external electrical field to straighten out the bent beam. Thomson also carried out deflection measurements, in which he applied a magnetic field to deflect the beam along with an external electrical field to straighten out the bent beam. From his measurements, he was able to calculate the charge/mass ratio of the electron: From his measurements, he was able to calculate the charge/mass ratio of the electron: e/m = -1.76x10 8 coulombs/gram

Sub-Atomic Particles Robert Millikan ( ) published the results of his Oil Drop Experiment in He designed an apparatus that could be used to determine the charge on an electron. Robert Millikan ( ) published the results of his Oil Drop Experiment in He designed an apparatus that could be used to determine the charge on an electron. The device used a fine mist of oil drops that had been exposed to ionizing radiation. The radiation caused some of the oil drops to take on one or more electrons. The device used a fine mist of oil drops that had been exposed to ionizing radiation. The radiation caused some of the oil drops to take on one or more electrons.

Sub-Atomic Particles

The Charge of the Electron

Sub-Atomic Particles Millikin determined that the charge on the electron is x coulombs. Millikin determined that the charge on the electron is x coulombs. Using Thomson’s value for the charge to mass ratio of the electron, the mass of the electron could be calculated. Using Thomson’s value for the charge to mass ratio of the electron, the mass of the electron could be calculated. mass of e - = (-1.60 x coulombs) (-1.76 x 10 8 coulombs/gram) (-1.76 x 10 8 coulombs/gram) = 9.11 x grams = 9.11 x kilograms = 9.11 x kilograms

Early Atomic Models J. J. Thomson had shown that all atoms contain negatively charged particles called electrons. Combined with the work of Millikan, they discovered that the electron has very little mass. J. J. Thomson had shown that all atoms contain negatively charged particles called electrons. Combined with the work of Millikan, they discovered that the electron has very little mass. Thomson proposed that the bulk of the atom is a positively charged gel or cloud, with most of the atomic mass and all of the positive charge uniformly distributed throughout the gel. Thomson proposed that the bulk of the atom is a positively charged gel or cloud, with most of the atomic mass and all of the positive charge uniformly distributed throughout the gel.

Early Atomic Models The electrons were viewed as discrete, very small particles that were stuck into the positively charged gel or cloud “like raisins in a pudding.” This model is often called the plum or raisin pudding model of the atom. The electrons were viewed as discrete, very small particles that were stuck into the positively charged gel or cloud “like raisins in a pudding.” This model is often called the plum or raisin pudding model of the atom. The electrons could be knocked out of the gel if enough energy is applied, and this is the source of the cathode rays. The electrons could be knocked out of the gel if enough energy is applied, and this is the source of the cathode rays.

Early Atomic Models One of the key features of Thomson’s atomic model is that most of the atomic mass and all of the positive charge is uniformly distributed throughout the atom.

Early Atomic Models Thomson had a graduate student, Ernest Rutherford, working for him. In 1911, Rutherford, Geiger and Marsden performed an experiment to confirm Thomson’s atomic model. Thomson had a graduate student, Ernest Rutherford, working for him. In 1911, Rutherford, Geiger and Marsden performed an experiment to confirm Thomson’s atomic model. They bombarded a thin gold foil with alpha (α) particles. The α particles have twice the charge of an electron and are positive in charge, with a mass that is 7300 times greater than the mass of an electron. They bombarded a thin gold foil with alpha (α) particles. The α particles have twice the charge of an electron and are positive in charge, with a mass that is 7300 times greater than the mass of an electron.

Early Atomic Models The α particles can best be thought of as a positively charged, fast traveling atomic sized bullet. They created a thin beam of α particles and directed the beam at a very thin gold foil. The α particles can best be thought of as a positively charged, fast traveling atomic sized bullet. They created a thin beam of α particles and directed the beam at a very thin gold foil.

Early Atomic Models If Thomson’s model is correct, most of the α particles should pass right through the gold atoms. Some slight deflection might occur if the positively charged α particle travels near an electron.

Gold Foil Experiment

Early Atomic Models The film that lined the apparatus showed that most α particles went through the foil with little or no deflection. However, some of the particles were deflected at great angles. The film that lined the apparatus showed that most α particles went through the foil with little or no deflection. However, some of the particles were deflected at great angles.

Early Atomic Models The deflection of the α particles was consistent with a large concentration of positive charge and atomic mass. This very small extremely dense positively charged area is called the nucleus.

Early Atomic Models The atom is mostly empty space, with the electrons found outside of the nucleus. If the nucleus was the size of a pea, it would have a mass of 250 million tons, and the electrons would occupy a volume approximately the size of a stadium.

Atomic Nucleus

Sub-Atomic Particles We now know that the positive charge of an atom, contained in the nucleus, is due to particles called protons. We now know that the positive charge of an atom, contained in the nucleus, is due to particles called protons. Protons have a charge equal in magnitude to that of an electron, but positive in charge. Protons have a charge equal in magnitude to that of an electron, but positive in charge. The mass of a proton is roughly 1800 times greater than the mass of an electron. The mass of a proton is roughly 1800 times greater than the mass of an electron.

Sub-Atomic Particles The nuclei of atoms also can contain neutrons. Neutrons are neutral in charge, with a mass similar to that of a proton. The nuclei of atoms also can contain neutrons. Neutrons are neutral in charge, with a mass similar to that of a proton. Neutrons are found in the nucleus of atoms, along with protons. Neutrons are found in the nucleus of atoms, along with protons.

Sub-Atomic Particles

During chemical reactions, atoms may lose or gain electrons to form charged particles called ions. During chemical reactions, atoms may lose or gain electrons to form charged particles called ions. Atoms of a given element may have differing numbers of neutrons. These forms of the same element are called isotopes. Atoms of a given element may have differing numbers of neutrons. These forms of the same element are called isotopes. It is the number of protons or the atomic number that defines the identity of the atom. It is the number of protons or the atomic number that defines the identity of the atom.

Atomic Symbols The periodic table lists the elements in order of increasing atomic number (the number of protons). The periodic table lists the elements in order of increasing atomic number (the number of protons). The atomic number, represented by the letter Z, is linked with the atomic symbol. For example, oxygen is atomic number 8, and any atom containing 8 protons, regardless of the number of neutrons or electrons, is represented by the symbol O. The atomic number, represented by the letter Z, is linked with the atomic symbol. For example, oxygen is atomic number 8, and any atom containing 8 protons, regardless of the number of neutrons or electrons, is represented by the symbol O.

Atomic Symbols To indicate a specific isotope, the atomic symbol must also contain the mass number. To indicate a specific isotope, the atomic symbol must also contain the mass number. The mass number is the number of neutrons plus protons for a particular isotope. The mass number is never found on the periodic table. The mass number is the number of neutrons plus protons for a particular isotope. The mass number is never found on the periodic table. Since the mass number is the number of particles (neutrons + protons) in the nucleus, it is always an integer. Since the mass number is the number of particles (neutrons + protons) in the nucleus, it is always an integer.

Isotopes of Sodium Mass number Atomic number

Atomic Symbols For example, there are three isotopes of carbon: For example, there are three isotopes of carbon: 12 C, 13 C and 14 C The mass number, if specified, appears in the upper left corner of an atomic symbol. Since all carbon atoms have 6 protons (carbon is atomic number 6 on the periodic table), atoms of carbon may have 6, 7 or 8 neutrons in the nucleus. The mass number, if specified, appears in the upper left corner of an atomic symbol. Since all carbon atoms have 6 protons (carbon is atomic number 6 on the periodic table), atoms of carbon may have 6, 7 or 8 neutrons in the nucleus. The isotopes are called carbon-12, carbon-13 and carbon-14.

Atomic Symbols If the atom has lost or gained electrons, the charge is written in the upper right corner of the atomic symbol. If the atom has lost or gained electrons, the charge is written in the upper right corner of the atomic symbol. The atomic number, though optional, may be written in the lower left corner of the symbol. The atomic number, though optional, may be written in the lower left corner of the symbol. 37 Cl 1- This ion of chlorine contains 17 protons, 20 neutrons, and 18 electrons.

Relative Atomic Masses Once the formulas of simple gases and compounds could be determined, scientists could also determine the relative masses of the elements. Once the formulas of simple gases and compounds could be determined, scientists could also determine the relative masses of the elements. For example, since equal volumes of gases contain equal numbers of particles (at the same T and P), the masses of gases could be compared to hydrogen, the lightest gas. For example, since equal volumes of gases contain equal numbers of particles (at the same T and P), the masses of gases could be compared to hydrogen, the lightest gas.

Stoichiometry Stoichiometry is a Greek word that means using chemical reactions to calculate the amount of reactants needed and the amount of products formed. Stoichiometry is a Greek word that means using chemical reactions to calculate the amount of reactants needed and the amount of products formed. Amounts are typically calculated in grams (or kg), but there are other ways to specify the quantities of matter involved in a reaction. Amounts are typically calculated in grams (or kg), but there are other ways to specify the quantities of matter involved in a reaction.

Relative Atomic Masses As the early chemists explored the nature of matter, they discovered that atoms of the elements had different masses. As the early chemists explored the nature of matter, they discovered that atoms of the elements had different masses. Avogadro’s Hypothesis which states that under the constant temperature and pressure equal volumes of gases contain an equal number of particles could be used to determine relative atomic masses for gaseous elements. Avogadro’s Hypothesis which states that under the constant temperature and pressure equal volumes of gases contain an equal number of particles could be used to determine relative atomic masses for gaseous elements.

Relative Atomic Masses Equal volumes of gases contain an equal number of particles. Equal volumes of gases contain an equal number of particles. Although the number of particles (atoms or molecules) in a liter of gas (at a specific T and P) wasn’t known, Avogadro’s Hypothesis said that a liter of any other gas under the same conditions would contain the same number of particles. Although the number of particles (atoms or molecules) in a liter of gas (at a specific T and P) wasn’t known, Avogadro’s Hypothesis said that a liter of any other gas under the same conditions would contain the same number of particles.

Relative Atomic Masses Equal volumes of gases contain an equal number of particles. Equal volumes of gases contain an equal number of particles. Since the masses of the gaseous samples could be determined, a comparative or relative scale of atomic and molecular masses could be derived. Since the masses of the gaseous samples could be determined, a comparative or relative scale of atomic and molecular masses could be derived.

Relative Atomic Masses Equal volumes of gases contain an equal number of particles. Equal volumes of gases contain an equal number of particles. For example, if the masses of a liter of oxygen (O 2 ), chlorine (Cl 2 ) and hydrogen (H 2 ) were compared under identical conditions, the hydrogen sample has the smallest mass, and the chlorine sample has the largest mass. For example, if the masses of a liter of oxygen (O 2 ), chlorine (Cl 2 ) and hydrogen (H 2 ) were compared under identical conditions, the hydrogen sample has the smallest mass, and the chlorine sample has the largest mass. The ratio of the masses of the 1 liter samples is: The ratio of the masses of the 1 liter samples is:35.5/16.0/1.00 Cl 2 / O 2 / H 2

Relative Atomic Masses Equal volumes of gases contain an equal number of particles. Equal volumes of gases contain an equal number of particles. The ratio of the masses of the 1 liter samples is: The ratio of the masses of the 1 liter samples is:35.5/16.0/1.00 Cl 2 / O 2 / H 2 Since all three gases are diatomic, we can say that an oxygen atom is 16.0 times heavier than a hydrogen atom, and that a chorine atom is 35.5 times heavier than a hydrogen atom. Since all three gases are diatomic, we can say that an oxygen atom is 16.0 times heavier than a hydrogen atom, and that a chorine atom is 35.5 times heavier than a hydrogen atom.

Relative Atomic Masses A scale of relative atomic mass was devised. Individual atoms are much too small to weigh, but the masses of large collections of atoms could easily be compared. A scale of relative atomic mass was devised. Individual atoms are much too small to weigh, but the masses of large collections of atoms could easily be compared. The relative masses of the atoms are listed on the periodic table. An arbitrary unit, the atomic mass unit (amu) is used for relative masses. The relative masses of the atoms are listed on the periodic table. An arbitrary unit, the atomic mass unit (amu) is used for relative masses.

Relative Atomic Masses Eventually, the carbon-12 isotope ( 12 C) was assigned an atomic mass of exactly 12 atomic mass units, and all other atomic masses are expressed relative to this assignment. Eventually, the carbon-12 isotope ( 12 C) was assigned an atomic mass of exactly 12 atomic mass units, and all other atomic masses are expressed relative to this assignment. The atomic mass for carbon, found on the periodic table, is amu, and not amu. This is because the periodic table lists the average relative atomic mass for all isotopes of the element. The atomic mass for carbon, found on the periodic table, is amu, and not amu. This is because the periodic table lists the average relative atomic mass for all isotopes of the element.

Relative Atomic Masses Carbon exists as three isotopes: Carbon exists as three isotopes: 12 C has a relative mass of exactly 12 amu 13 C has a relative atomic mass of amu 14 C has a relative atomic mass of 14.0 amu The value found on the periodic table, amu, reflects the relative abundance of the isotopes. The majority of carbon (98.89%) is The value found on the periodic table, amu, reflects the relative abundance of the isotopes. The majority of carbon (98.89%) is 12 C, with 1.11% 13 C, and a trace of 14 C. 12 C, with 1.11% 13 C, and a trace of 14 C.

Relative Atomic Masses Chorine exists as two isotopes: Chorine exists as two isotopes: 35 Cl, with a relative atomic mass of 35.0 amu and 37 Cl, with a relative atomic mass of 37.0 amu. What does the atomic mass of chlorine on the periodic table tell you about the relative abundance of the two isotopes?

Moles Many chemical reactions are carried out using a few grams of each reactant. Such quantities contain huge numbers (on the order of ) of atoms or molecules. Many chemical reactions are carried out using a few grams of each reactant. Such quantities contain huge numbers (on the order of ) of atoms or molecules. A unit of quantity of matter, the mole, was established. A mole is defined as the number of carbon atoms in exactly 12 grams of 12 C. A unit of quantity of matter, the mole, was established. A mole is defined as the number of carbon atoms in exactly 12 grams of 12 C. Avogadro determined the number of particles (atoms or molecules) in a mole. Avogadro determined the number of particles (atoms or molecules) in a mole.

Moles Avogadro’s number = x particles/mole Avogadro’s number = x particles/mole Atoms are so small, that a mole of most substances can be easily held in ones hand. Atoms are so small, that a mole of most substances can be easily held in ones hand. Cu Al Hg Fe I2I2 S

Moles If we consider objects we can see, a mole of sand would cover the entire planet and be several miles deep! However, the collection of atoms, called a mole, is very convenient in the laboratory (just like dozens are useful in buying eggs or pencils). If we consider objects we can see, a mole of sand would cover the entire planet and be several miles deep! However, the collection of atoms, called a mole, is very convenient in the laboratory (just like dozens are useful in buying eggs or pencils).

Moles A mole of any atom has a mass equal to the element’s atomic mass expressed in grams. A mole of any atom has a mass equal to the element’s atomic mass expressed in grams. A mole of iron atoms has a mass of grams; a mole of iodine molecules (I 2 ) has a mass of (126.9) (2) = grams.

Moles The masses of each molar sample are provided below. The masses of each molar sample are provided below. Cu = 63.55g Al=26.98g Hg = g Fe=55.85 g I 2 =253.8 g S=32.07g

Molar Mass For compounds, once the formula is known, the mass of a mole of the substance can be calculated by summing up the masses of all the atoms in the compound. For compounds, once the formula is known, the mass of a mole of the substance can be calculated by summing up the masses of all the atoms in the compound. For example, hydrogen peroxide has the formula H 2 O 2 : For example, hydrogen peroxide has the formula H 2 O 2 : 2H +2O = 2(1.008g) + 2(16.00g) = g/mol A mole of hydrogen peroxide has a mass of grams. A mole of hydrogen peroxide has a mass of grams.

Nuclear Chemistry Many nuclei are unstable, and may emit particles and/or energy until they reach a more stable combination of protons and neutrons. A nuclear reaction differs from a chemical reaction, and results in the formation of new elements as an unstable nucleus decays.

Nuclear Stability For the lighter elements in the periodic table, most of the stable nuclei contain approximately an equal number of protons and neutrons. For heavier elements, the neutron to proton ratio generally increases in the stable isotopes.

Nuclear Stability

The mass of a nucleus is less than the sum of the masses of its neutrons and protons. A very small amount of mass is converted into the nuclear binding energy. During nuclear reactions, some of this energy is released. A very small loss of mass results in a huge release of energy.

Nuclear Binding Energy

Nuclear Stability Smaller atoms will undergo fusion, and combine to form heavier atoms. Larger atoms will undergo fission, and break into smaller nuclei and particles.

Nuclear Equations Nuclear equations represent only the nuclei of the particles, rather than the atom complete with electrons. Nuclear equations must be balanced. This means that the sum of the mass numbers on the left side equals the sum of the mass numbers on the right side. Nuclear equations must be balanced. This means that the sum of the mass numbers on the left side equals the sum of the mass numbers on the right side.

Nuclear Equations Likewise, the sum of the atomic numbers on the left must equal the sum of the atomic numbers on the right. Note that the identity of the elements will change during a nuclear reaction.

Nuclear Reactions Unstable nuclei can decay in a variety of ways. If the nucleus contains too many protons/neutron, βemission may occur. A β particle is an electron, but it results from a neutron disintegrating into a proton and an electron. The proton remains in the nucleus of the newly formed atom, and the electron is ejected.

βEmission The net result of β emission is an increase in the number of protons in the nucleus.

βEmission Nuclear reactions must be balanced in terms of the total of the atomic numbers and the total of the mass numbers.

α Emission Some unstable nuclei emit an α particle. α particles are the same as a helium nucleus. The alpha particle contain two protons and two neutrons. The net result is a decrease in both atomic number (by 2) and mass number (by 4).

α Emission Alpha emission is common in heavier elements. All elements above 209 Bi are radioactive. They often emit alpha particles (along with other forms of radiation) until they attain a stable nucleus.

α Emission

γ Emission Gamma (γ) emission is emission of a very high energy form of radiation. It accompanies both α and β emission, and represents a release of energy. Since it is not a particle, it is often omitted from nuclear equations.

Positron Emission Nuclei with a neutron to proton ratio which is too low, may undergo positron emission. A positron, β +, has the same mass as an electron, but it is positive in charge. The positron results from a proton decaying into a neutron and ejecting the positron. As a result, the number of neutrons increases, and the number of protons decreases. The positron results from a proton decaying into a neutron and ejecting the positron. As a result, the number of neutrons increases, and the number of protons decreases.

Positron Emission A positron has the same mass as an electron but an opposite charge. It can be thought of as a “positive electron.” Positron,  +

Electron Capture Nuclei with a neutron to proton ratio which is too low, may undergo electron capture. In this process, a proton in the nucleus combines with an electron outside the nucleus to form a neutron. The net result is an decrease in atomic number and an increase in the mass number. The net result is an decrease in atomic number and an increase in the mass number.

Electron Capture