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Atomic Energy Levels - Video Clips

Evolutionary Patterns - Video Clips


Atoms and molecular structure


Chemistry is the study of matter and the reactions of matter. The use of chemistry creates new types of glass, plastics, paints, fabrics, drugs, flavors, metals, and many other materials. The molecular secrets of DNA, ceramics, lasers, medicines, cosmetics, etc. are areas of chemical study. And chemists continue to improve reactions that lead to fertilizers, rocket fuel, air bags, fuel cells, and more. Your study of chemistry is one that will allow you to understand the vocabulary that chemists use to identify different types of materials and reactions. You will also learn the reasons for different properties of matter such as boiling point and solubility. And, of course, you will study types of reactions from simple decomposition reactions to acid and base reactions.

Essential vocabulary

In previous science classes you were introduced to different classifications of matter and examined different properties of matter. The vocabulary from these earlier classes is necessary to explain and describe chemistry. Thus it is essential that you know definitions of the following words fully.

Matter is something that takes up space and has mass. All substances are matter. Matter is not energy, like heat or light, and it is not a force, like electrical force, magnetic force, or a push or pull.

Substance is a term used to identify a sample of matter that is being considered. A pure substance is either an element or a compound and is one type of matter with one uniform set of properties. Substances can either be pure substances or mixtures.

Atoms are the individual particles that make up the simplest forms of matter, which are elements. If you break apart an element over and over again you will finally get to the smallest particle that is still the same element: an atom. Atoms do have other parts and can be broken, but if you break an atom the pieces are not the same as the element you started with.

Elements are the simplest form of matter. If you have a pure form of an element then you cannot break it down by any chemical reaction to make a new substance. An element is composed of a single type of atom.

Compounds are pure substances with one type of properties (that is they are not mixtures). While compounds are single types of matter, like water, carbon dioxide, sugar, and salt, they can be broken down into elements or simpler compounds with a chemical reaction. For example, electricity can be used to separate water, H2O, into the elements hydrogen, H2, and oxygen, O2.

Molecule is the smallest particle of a pure substance that still has the properties of the substance. A molecule of water would be one H2O, an oxygen with two hydrogens held together by chemical bonds. A molecule of hydrogen, H2, is two hydrogen atoms combined with a chemical bond.

Mixtures are combinations of different types of substances. Sugar water, paint, milk, air, brass, granola bars are examples of mixtures. Heterogeneous mixtures are chunky and look like mixtures with different, recognizable parts. Salads and chocolate chip granola bars are good examples. Homogeneous mixtures can easily be mistaken for compounds because they look like they have only one set of properties. But if a substance can be separated without a chemical change then it is a mixture. For example, sugar water can be separated by letting the water evaporated or air can be separated by using special filters and temperature changes that separate the nitrogen, oxygen, water, carbon dioxide, and argon.

A common way of showing the difference and similarities between the different types of matter is to use a flow chart.

Describing the Properties of Matter

When an experiment is conducted it is necessary to describe the substance and the changes the substance undergoes during the experiment. There are two types of changes and properties: physical and chemical.

A physical change is a change that takes place without creating new chemical substances. Filtering, evaporation, melting, floating (or sinking), separating by color or size, and distillation (a technique that uses boiling point differences) are some methods used for separating mixtures.

A chemical change is a change that creates a new substance. A chemical change is the same as a chemical reaction. Burning, fermenting or rotting, decomposing, and synthesizing or combining chemicals are types of chemical changes. A chemical change can be identified by a change in physical properties like color, odor, state of matter, or texture; or a change in energy, which may be absorbed or released.

Properties are descriptions of substances and can either be physical or chemical. Physical properties are determined without creating a new substance: boiling point, color, texture, shape, odor, density, etc. Chemical properties are determined by creating new substances and are generally a brief description of a chemical reaction: burns rapidly, ferments in warm conditions, reacts with acids to produce hydrogen gas, is explosive when lit.

All the bold and italicized terms are important vocabulary for chemistry. Please refer back to these definitions whenever you need a remind yourself of the meanings of terms or better yet develop a set of flash cards for study and reference.

Atomic and Molecular Structure

  • The First Periodic Table


The basis for the science of chemistry was the discovery and classification of elements. The first elements were known at the beginning of civilization. Iron tool, sulfur medicinal balms, gold jewelry, and carbon paints were all used in many ancient civilizations. When chemists started looking for elements some of them were easy to find. Yttrium, ytterbium, erbium, and terbium were all named by the Swedish research groups for the town of Ytterby, Sweden where the ore containing the new elements was mined. The elements of salt, sodium and chlorine where identified in 1807 and 1774 respectively. Cobalt was discovered in 1735, and oxygen in 1774. Aluminum was purified and identified in 1825 and phosphorous in 1669. Some elements have been difficult to isolate. Technium, in the middle of the periodic table, was finally characterized in 1937. Francium continues to be elusive but was purified in sufficient amounts to identify in 1939. And astatine at the bottom of the halogens was characterized from a small radioactive sample in 1940. We have been creating and identifying elements from 1945 when neptunium was created in California to present when a joint Russian and American group isolated 6 atoms of element 117 in 2010 (the last named element was copernicium, Cn, is element #112).

Many chemists attempted to develop useful lists of elements before the modern periodic table was developed by Dmitri Mendelev. First published by Mendelev in 1908, the periodic table lists all the elements and some of their properties. Mendelev knew that elements have similar physical properties and reactions with common reactive nonmetals, chlorine and oxygen. Using chemical properties and the physical properties of elements such as melting point, he organized elements into families, which are elements that have similar properties. Then placing the families in columns he arranged the families so the atomic mass increased from lightest to heaviest. The result was the first periodic table, which remains the backbone of today’s periodic table.

Mendeleev’s Periodic Table

Alkali Metals Alkaline Earth Metals Halogens Noble Gases Transition Metals Lanthanides Actinides Semimetals , , , , , All are Metals , , All are Nonmetals The Modern Periodic Table & Group Names

The periodic table is a list of all elements organized so that the atomic number increases from lowest number to highest number. The first 92 elements are created by natural events and the other elements (26 more at this time) are made by man. The list has rows called periods that stop when the next element has properties similar to the properties of the first column of elements. Each column is a group or a family.

Periodic Groups There are two broad groupings of elements that are separated on the periodic table. These are the metals and the nonmetals. Metals are on the left and extend from lithium, Li, down to francium, Fr, then across to the right to aluminum, Al, (not borium, B) to bismuth, Bi. The nonmetals are carbon, C, and the diagonal line to radon, Rn and all elements to the right of that diagonal line. Hydrogen is a nonmetal. This leaves a set of elements that have properties that are not fully metallic called semimetals (or metalloids): boron, B; silicon, Si; gallium, Ga; arsenic, As; antimony, Sb; tellurium, Te; and polonium, Po. Semimetals slightly conduct electricity and may be brittle and without luster.

Four groups on the periodic table have distinct names and distinct properties. The groups alkali metals, alkaline earth metals, halogens, and noble gases. When we understand the chemistry of one element in any of these groups then we know that the other members of the group have similar reactions and properties. (Recall groups are columns of elements.)

The first column of the periodic table is the alkali metals: lithium, Li; sodium, Na; potassium, K; rubidium, Rb; cesium, Cs; and francium, Fr. Alkali metals are soft metals that can be easily cut, have low melting points, have low densities (lithium, sodium, and potassium float on water), react with oxygen and water quickly. They form oxygen compounds in a 2 to 1 ratio, e.g., Na2O and K2O, and form chlorine compounds in a 1 to 1 ratios, e.g., NaCl and KCl (great video at

Alkali metals: lithium, sodium, potassium, rubidium, and cesium ( ; ; ; ; )

Note that hydrogen, H, is often placed in the same column but it is not an alkali metal, in fact, hydrogen is unique and is alone in its own family.

The second column of the periodic table is the alkaline earth metals: beryllium, Be; magnesium, Mg; calcium, Ca; strontium, Sr; barium, Ba; and radium, Ra. These metals have similar properties to alkali metals except they are harder, have higher melting points and densities, and react more slowly with oxygen and water. The most important change is the change in the ratio of compounds formed with oxygen and chlorine. With oxygen these metals form compounds with a 1 to 1 ratio, like MgO and CaO, and react with chlorine to form compounds with a 1 to 2 ratio, like MgCl2 and CaCl2.

The second to last column is the halogens: fluorine, F; chlorine, Cl; bromine, Br; iodine, I; and astatine, At. These are toxic, dangerous, and reactive substances (isolation of fluorine took the lives of several early experimenters). They are so reactive they are found in nature combined into a diatomic molecule such as F2, Cl2, Br2, and I2 (At is radioactive and is only characterized as atoms). Halogens are the only elements with -ine endings so they are easy to recognize). Halogens easily form gases. Fluorine and chlorine are gases at room temperature and bromine is a liquid (the only other liquid at room temperature is mercury, Hg). Iodine so easily forms a gas that even though it is a solid it will sublime (change from solid to gas) at room temperature. These elements are also commonly found in salts such as common table salt, NaCl, and ocean salts, KCl and MgCl2.

The last column is the noble gases: helium, He; neon, Ne; argon, Ar; krypton, Kr; xenon, Xe; and Radon, Rd. These are all gases nonreactive, hence their name referring to being too good to react. Recently, under unique conditions Kr, Xe, and Rd have formed stable compounds. The fact that they are gases and are not reactive meant that they were very hard to discover and isolate.

Noble Gases excited by Electricity (

Some other group names are the transition metals which are the elements in the “middle” of the periodic table starting with scandium, Sc, (#21) across to zinc, Zn, and down to mercury, Hg, and Element 112 (copernicium).

Also two periods of elements separated from the other elements in most periodic tables are the lanthanoids and actinoids (commonly referred to as the lanthanides and actinides) that start at lanthium, La, (#57) . While these elements are commonly referred to as rare earth metals, they are actually rather common elements. Also Glenn Seaborg (who created several transuranic elements like plutonium and americium) suggested, much to the distress of his elders, that the family names be given for the periods not the groups, because he rightly pointed out that the properties of these elements were similar in the rows not the columns.

Numbers on the Periodic Table

The periodic table is organized by increasing atomic number. Each element is different from its neighbor by the number of protons present in the nucleus. (We will learn about protons and neutrons in the Fundamental Particles section of the text). The atomic number is both an identifier for the element and also a count of the number of protons present in the nucleus. The second number commonly found on the periodic table is the atomic mass. This number is much easier to determine experimentally than the atomic number so it was originally used to determine the order of the elements. The atomic mass is an average mass of all the atoms in a sample and is essentially the sum of the two types of particles in the nucleus: protons and neutrons. (This last statement should be troublesome. How can iron, Fe, have particles that add up to 55.85 particles? This has to do with how the average of masses are determined. When we get to Isotopes this will be dealt with in more detail.)

It turns out that ordering the elements by their atomic number creates a periodic table that is better at predicting properties, chemical reactivity, and electron configuration, than one based on atomic mass. So there are four main element pairs where the atomic number increases but the atomic mass decreases: Ar and K (18 to 19 but 39.9 to 39.1), Co and Ni (27 to 28 but 58.9 to 58.7), Te and I (52 to 53 but 127.6 to 126.9), and Th and Pa (90 to 91 but 232.0 to 231.0).

A Final Word about the Periodic Table The periodic table is one of the most valuable tools in chemistry; in fact, there is another section coming up soon on periodic trends. The fact that the elements that make up all things can be placed in an organized pattern has helped scientists make sense of the chemistry and properties of substances. Furthermore, the organization of the table is based on the nature of the atom and the particles that make it up. As we go through the different topics use the periodic table to recognize patterns in properties and reactions to make your study of chemistry more efficient and complete.

The Atomic Model

The modern model of the atom is built from bits and pieces that scientists have added to the study of chemistry. Important names like Thompson, discoverer of the electron, Chadwick, discoverer of the neutron, Millikan, discoverer of the charge of the electron, and Mosely, who figured out how to determined the atomic number in 1903, are given this brief mention though their contributions were enormous. There are many others and you are encouraged to discover for yourself the rich history behind the present day atom. So while the atomic model was developed after a series of hypothesis and theories were put forward, reworded, replaced as new experiments required new theories. What we know today as certainty, may change if new tools are developed that are sensitive enough to make the keen observations necessary to fully support new theories. (see and listen to

Gold Foil Experiment The one experiment that provided the strongest evidence for the nuclear model of the atom is Rutherford’s Gold Foil Experiment. At Rutherford’s labs in Cambridge, England, they set up an experiment where a radioactive particle, an alpha particle, is fired at a very thin sheet of gold. The expectation is that the particles will fly though the thin gold foil, because it is thought that gold’s mass is distributed throughout its volume (the alpha particle is dense and small like a bullet and the foil is like paper). But much to the astonishment of the researchers some particles are reflected back away from the foil while most make their way through the foil. This scenario made it impossible to use a model where the atom was a solid sphere like a marble. After much contemplation (and repeating the experiment), Rutherford theorized that the atom must have a small, dense, and positive nucleus. This would explain the small number of particles that bounced back, 1 in 8800, and why the nucleus wouldn’t fall apart with the collisions. This was the beginning of our understanding of the modern, nuclear atom.

The modern atom has a small dense nucleus in the center of a sphere. This nucleus is 10,000 times smaller than the outer boundary of the atom, but it must contain nearly all the mass of the atom. This is an astounding notion. For example, imagine a two-story house that weighs 100,000 pounds, but instead of the mass being contained throughout the house it is all located in a pea dropped in the middle of the second floor. So you can lift up all the parts of the house because they have very little mass, but when you get to the pea it won’t budge, because it weighs as much as the entire house.

The volume around the nucleus, called the electron cloud, is filled with electrons that are moving near the speed of light (about two million meters per second– These electrons move randomly in different directions and different speeds. 90% of the time the electrons are within the atom’s radius. And like some super freeway the electrons are speeding so fast nothing will cross their path. This is how atoms have outer boundaries that can be measured.

Parts in the Atom and their Size

Fundamental Particles in the Atom The proton is located in the nucleus and has a mass that is the nearly identical to neutron’s and 1836 times bigger than an electron. The atomic number of an element is equal to the number of protons and differs for each element. A proton has a positive charge. The neutron is also in the nucleus and has a neutral charge. When the number of neutrons changes, the new atom is an isotope. Isotopes are atoms with the same number of protons but different number of neutrons. Another definition is that isotopes are atoms of the same element with different mass numbers. The electrons are negative particles with very little mass and travel throughout the nucleus’s volume. The number of electrons in a neutral atom is equal to the number of protons. The electrons have a negative charge with very little mass. Electrons are also responsible for the chemistry of substances because the nucleus is so deeply buried in the core of the atom.

Fundamental Particles in an Atom Name Location Relative Mass Charge electrons electron cloud 1/1836 negative, e– proton nucleus 1 positive, p+ neutron nucleus 1 neutral, n0

Quarks Very recently scientists discovered smaller pieces of protons and neutrons called quarks (1968 to 1995— Quarks and electrons seem to be the smallest particle that makes up matter. Three quarks are needed to make up both a proton and a neutron, but a proton is made up of two “up” quarks and one down quark so that the total charge is +1, and the neutron has two down quarks and 1 up quark so that the charge is neutral (up quarks have a charge of +2/3 and a down quark has a charge of –1/3).


Charge is a unique property that exerts a force between charged matter. If matter has the same charge then there is a pushing force between the objects. If the substances have opposite charge then the substances are attracted (“like charge repels, unlike charges attract”). But since protons are stuck in the nucleus and cannot move (changing protons changes the element), positive matter is created when electrons are lost and negative matter is created when electrons are gained. The static electricity of a balloon caused by rubbing it against a fuzzy shirt or hair is a property of charge. The balloon gains electrons from the fuzzy clothing and has a negative charge. The shirt or hair is positively charged with the lost electrons. The oppositely charged balloon and hair stick to each other with an attractive force. This is a very strong force, since we move only relatively few electrons and yet overcome the force of gravity. Charge is important in understanding the bonding taking place within and between elements, as well as, the interaction between electrons and between electrons and the nucleus.

Balloons and hair show electrons move to create static charge

Determining the Number of Fundamental Particles in a Neutral Atom The number of protons in any atom of an element is equal to the atomic number. The number of electrons is equal to the number of protons. This makes the atom neutral, which is when the sum of positive and negative charges cancels out.

Atomic Mass and Neutrons The atomic mass of an element is the combination of the protons and neutrons found in the nucleus of the atom. For fluorine with an atomic number of 9 and atomic mass of 19.0 amu, the number of protons is 9 and the number of neutrons is 10 (since 19 = 9 + 10). For chromium the atomic number is 24 and the atomic mass is 52.0 amu (amu is the unit for mass for an atom that means atomic mass unit). So the number of protons for chromium is 24 and the number of neutrons is 28 (since 52 = 24 + 28).

But how does this work for something like boron, B, with an atomic mass of 10.80 or for chlorine, Cl, with an atomic mass of 35.45. Atomic mass is actually a calculated number based on the different isotopes of an element. Isotopes are atoms of the same element with different number of neutrons. For example, boron has two major isotopes,, boron-10 and boron-11. Boron-10 has 5 protons and 5 neutrons and boron-11 has 5 protons and 5 neutrons. The 10 in boron-10 is called the mass number. Since fluorine and chromium a single isotope that is 100% of the atoms, then the calculation is not needed.

The calculation to determine atomic mass is called a weighted average. For example, boron has two major isotopes, boron-10 and boron-11, and all natural samples of boron contain 20% boron-10 and 80% boron--11. The average mass of the atoms in a natural sample is calculated by: 0.20 • 10 + 0.80 • 11 = 10.8 amu (a more exact calculation of mass would use the actual mass of the isotope). Here is another example using chlorine. Chlorine also has two major isotopes: chlorine-35 and chlorine-37 with natural abundances of 75% and 25% (natural abundance is the percent of an isotope found in nature). So this weighted average calculation is: 0.75 • 35 + 0.25 • 37 = 35.5.

Electron Configuration and the Periodic Trends Our model of the atom has the electrons on the outer boundary and in this location they will be the particles that interacts with other substances in chemical reactions. But we also know from the groups on the periodic table (alkali metals, alkaline earth metals, halogens, and noble gases) that the same reactivity and properties reoccur as more electrons are added. This means that adding electrons (recall electrons equal the protons in a neutral atom) does not gradually change the properties of atoms, but rather properties of atoms repeat as electrons increase. Therefore, the arrangement of electrons changes across the periodic table (to change properties) and then the arrangement returns to something similar so that the properties of elements repeat in the families of elements. Thus, the origin of an elements chemistry must lie in the arrangement of electrons.

The Bohr Model of Electron Configuration. The arrangement of electrons in the atom also needs a theory that predicts and explains the properties. Many of the early theories suggested that the electrons traveled like planets in orbits about the nucleus. The problem with this theory was that it was known that an object traveling around an attractive center will slowly move towards the center. So if the orbit theory was true then some atoms would be spontaneously changing as electrons and protons reacted when the electron spiraled into the attractive, positive nucleus.

Below is a brief description of the quantum model of the atom. You are encouraged to look at these two sources for a complete description: Physics 2000 at and Hyperphysics at .

Neils Bohr suggested that electrons had to remain in specific energy levels that had fixed energies. This restriction forced electrons to maintain their energy. In support of this theory, he showed calculations of the emission spectra of hydrogen gas that matched his theory. Emission spectra is the release of light after energy has been added to a substance. Unfortunately, this was the only substance that Bohr’s model fit. Fortunately, Bohr was able to put together a number of ideas and theories and point the way to the model we use today.

Bohr Model and Explanation

Mathematics of Waves and Light. The mathematics needed for a deep understanding of quantum theory is quite complex. Here are two simple, basic equations that show the relationship between the energy of light emitted (or absorbed) by an atom and the energy difference between energy levels that the electrons travel in. Different types of light have different frequencies, ?. When you are changing radio channels you are changing frequencies that are emitted from each radio stations (light is all parts of the electromagnetic spectra which includes radio waves, microwaves, infrared waves, visible light, ultraviolet light, x-rays, and gamma rays). Another measurement of light is wavelength, ?. Frequency and wavelength are related to each other by the speed of light, c: c = ?•? (the speed of light is a constant through a vacuum, 3 X 108 m/s).

Energy of any frequency of light is calculated by using another constant, h. h is plank’s constant and equals 6.63 X 10–34 m2kg/s. Plank’s constant is the smallest step of energy that can separate energy levels in an atom or between frequencies of light (on a normal scale energy is continuous because Plank’s constant is extremely small and differences in the steps are unobservable). Thus E = h? and has units of measurement of joules, J.

Hydrogen Emission Spectrum (Visible Light, Balmer Series) (

These two equations were used by Bohr to support his model of the arrangement of electrons for the hydrogen electron. Energy is emitted as light when an electron loses energy to move to a lower level. But the change has to take place in steps of h, Plank’s constant. If the change was gradual or varied, then the bands of light would broaden and create a rainbow of light. Since the change has to be in steps, the lights appear as distinct lines. Each color is related to a wavelength, ?, or a frequency, ?. The frequency and wavelength of light is related to the energy, E = ?•h, needed to change energy levels. Comparing the observed light spectra with the energy predicted by theory, provides evidence for the correctness of the theory.

Schödinger and others showed that the behavior or electrons follows the pattern of waves better than as objects that obeying traditional physics like the movement of planets, which is what Bohr’s model was based on. Awkward sentence. They modified Bohr’s model using wave equations, or functions, to model the energy of electrons. The result is the quantum model, the Schödinger model, of the atom that correctly predicts the behavior of electrons and provides us with a physical representation of the location of electrons that we call electron configuration (electrons in compounds can also be modeled by quantum theory, but such models are much more complex).

Quantum Numbers The wave equations used to describe the electron configuration require at least 3 variables called quantum numbers. Changing the quantum numbers, the variables in the wave equation, changes the energy of the electron. Each electron has its own unique set of quantum numbers.

The shapes shown for the quantum orbital are created by graphing the wave equations in three dimensions. The outer boundary of the shape is the cut-off point for the location of the electron 90% of the time. Since the electron in an orbital travels randomly both in direction and speed, pinpointing its location or path is impossible (this is the famous Heisenburg uncertainty principle: “The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa.”). Thus an orbital represents the “container” where an electron is probably located, and each orbital is a graph of the wave equation for a specific energy.

The First Quantum Number: Energy level or shell, n The first quantum number describes the radius the electron exists away from the nucleus. So n = 1 means that the electrons with this quantum number will be closest to the nucleus. n = 2 will be in the second layer, n = 3 will be in the third layer, and so on up to n = 7.

The Second Quantum Number: Orbital Shape subshell, l The second quantum number determines the shape of the orbital, which is the location where electrons can be found in the atom. The number tells you how many nodes are in the shape, where a node is the crossing point in the shape. So l = 0 has no crossing points and will be just a sphere like we expect of an atom (recall the electron location, or electron cloud when all orbitals are combined, is the boundary of the atom). l = 1 has 1 crossing point and looks like a figure-8. l = 2 has 2 crossing points and looks like a four-leaf clover. l = 3 has 3 nodes and is a bit complicated (l = 4 is the highest orbital for this number). The shapes are labeled by either numbers or letters: l = 0 = s; l = 1 = p; l = 2 = d; l = 3 = f.

l = 0 = s l = 1 = p l = 2 = d l = 3 = f

The Third Quantum Number: Orientation in Three Dimensional Space The third quantum number, ml is used to designate orientation in space. The figure-8 shape with l = 1, has three shapes needed to completely fill the spherical shape of an electron cloud. That is combining figure-8 shapes going up and down (z-axis), left and right (x-axis), and back and forth (y-axis), the entire electron cloud will be filled with figure-8 orbitals. It takes three orientations for figure-8 orbital, l = 1, five for the cloverleaf orbital, l = 2, and seven for l = 3 ( l + 1, or consecutive odd numbers).

The 3 p-orbital

(s orbital has only one orientation, it changes only its size)

(The upper dots represent multiple “snapshots” showing the location of an electron over time.)

The 5 d-orbital The 7 f-orbitals

The Spin Quantum Number: Only Two Electrons per Orbital There is a fourth quantum number that doesn’t describe the shape or energy of an electron. This is the spin quantum number, ms. It turns out that quantum theory requires only two electrons to exist inside any single orbital, where a single orbital has its own size, shape, and orientation. So the two electrons have different spins with ½ spin (up) and –½ spin (down). Spin is a property an electron exhibits when it is placed in a magnetic field.

Certain values are allowed for the different quantum numbers: n, can be 1 to 7; l, can be (n – 1) to 0 for each energy level n; ml can be -l to 0 to l for each subshell, l; and spin is ½ then –½ for every orbital. This means that certain combinations of quantum numbers are not allowed. For example, the following quantum numbers break the rules: (0,0,0,½) because n ? 0; (2,2,-1, ½) because l must be smaller than n; (3,2,3,½) because ml is larger than l; (3,2,-3,½) because ml is smaller than -l; and (1,0,0, ¾) since ms must be only ½ or -½.

Summary of Quantum Numbers Variable n l ml ms Names of Quantum Numbers Principle Quantum Numbers Angular Momentum Quantum Number Magnetic Quantum Number Electronic Spin Quantum Number Possible numbers n = 1,2,.., 7 l = n – 1 to 0 s=0, p=1, d=2, f=3 ml = -l,...,0,...,l 1/2 or -1/2 Also Known As Energy Level or Shell Shape of the Orbital or Subshell, l equals the number of nodes or crossing points Orientation or Alignment along an Axis or in a Plane Spin Physical Representation Distance from the nucleus s-orbital is a sphere p-orbital is a figure–8 d-orbital is a cloverleaf f-orbital has 8 & 6 lobes For example, the figure-8 shape may lie on the x, y,&z-axes. d & f-shapes have planes of symmetry. Arrows up or down for magnetic alignment

Periodic Table Characteristic Row number Periodic table blocks 1/2 of each block is assigned a number line with 0 in the middle. The first half of a block has spin up and the second has spin down.

The following examples show how the quantum numbers are translated into three- dimensional shapes.

Examples of Quantum Numbers and Shapes Description Quantum Numbers 3D Shape First Energy Orbital, lowest energy (1, 0, 0, 1/2) n = 1 first energy level l = 0 = s-orbital shape (sphere) First Energy Orbital, but for 2nd electron (1, 0, 0, –1/2) Same orbital and energy as above. but second electron has opposite spin. (sphere) Second Energy Level Orbital (2, 0, 0, 1/2)

n = 2 = second energy level l = 0 = s-orbital (larger sphere) Second Energy Level p-orbital (2, 1, 0, 1/2)

n = 2 = second energy level l = 1 = p-orbital (higher energy than s–orbital) (figure-8) (ml = 0) Second Energy Level p-orbital (2, 1, 0, 1/2)

Same orbital and energy as above, but oriented on a different axis (three possible orientations). (figure-8) (new . orientation, ml = 1) Third Energy Level with d-orbital shape (3, 2, 0, 1/2)

n = 3 = third energy level l = 2 = d-orbital (five possible orientations) (larger 3rd orbital)


Using the Periodic Table to Determine Quantum Numbers Quantum numbers are unique to each electron. For example the quantum numbers for the first ten electrons are all different: 1st = {1,0,0,½} ; 2nd = {1,0,0,–½} 3rd = {2,0,0,½} ; 4th = {2,0,0,–½} ; 5th = {2,1,–1,½}, 6th = {2,1,0,½} ; 7th = {2,1,1,½} 8th = {2,1,–1,–½}, 9th = {2,1,0,–½} ; 10th = {2,1,1,–½}

For every atom these quantum numbers would be used for the first ten electrons, because they represent the lowest ten energies for electrons. Remarkably the periodic table, developed by examining the chemical and physical properties of elements, provides a clear diagram for determining the quantum numbers of electrons.

To use the periodic table to determine quantum numbers it must be divided into blocks and the blocks divided in half (and square 2 with He is moved next to square 1 with H.) The row numbers mostly match the principle quantum number, n (d-block and f-block must use n’s one less than the row number and two less than the row number respectively). The blocks match the angular momentum quantum number, l, and both letter and number values are given. ml is shown as a pair of number lines and ms is positive for the first half of the block and negative for the second half.

s = 0 p = 1

1 2

2 3 4 d = 2 5 6 7 8 9 10 3 11 12 13 14 15 16 17 18 4 19 20 * 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 5 37 38 * 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 6 55 56 * 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 7 87 88 * 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 ml 0 0 -2 -1 0 1 2 -2 -1 0 1 2 -1 0 1 -1 0 1 ms ½ -½ ½ ½ ½ ½ ½ -½ -½ -½ -½ -½ ½ ½ ½ -½ -½ -½

Here are some examples of quantum numbers of electrons using this divided periodic table: The 1st electron in any atom has {1,0,0,½}: row 1, s-block, ml =0, ms=0 The 8th electron in any atom has {2,1,-1,-½}: row 2, p-block, ml =-1, ms=-½ The 75th electron in any atom has {5,2,2,½} : row 6 but for the d-block the n, principle quantum number is one less than the row number, hence the 5, d-block, ml =2, ms=½ The 84th electron in any atom has {6,1,-1,-½}: 6th row, p-block, ml =-1, ms=-½

If we add the f block (with f = 3) 6 **57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 7 **89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 ml -3 -2 -1 0 1 2 3 -1 -3 -2 -1 0 1 2 3 ms ½ ½ ½ ½ ½ ½ ½ -½ -½ -½ -½ -½ -½ -½ -½

Using the divided periodic table for electrons in the f-block, the 68th electron’s quantum numbers are {4,3,0,-½}; 6th row but back 2 for the f-block so 4 is used, f-block = 3, ml =0, ms=-½ And the 94th electron has {5,3,2,½}.

A common question about quantum numbers might ask, “What is the highest energy electron for the element scandium, Sc?” To answer this question you would find the last electron needed for Sc, which is electron 21, and then use it’s quantum numbers: {3,2,-2,½}.

Just remember that electrons in the d-block have n equal to the row number – 1 (row number minus one) and the electrons in the f-block have n equal to the row number – 2.

Electron Configuration The purpose of introducing quantum numbers has been to show that similarities in the electron arrangement or electron configuration lead to the similarities and differences in the properties of elements. But writing the quantum numbers of electrons of an element in set notation like {2,1,-1,½} is time consuming and difficult to compare so an abbreviated form was developed. An electron configuration lists only the first two quantum numbers, n and l, and then shows how many electrons exist in each orbital. For example, write the electron configuration of scandium, Sc: 1s2 2s2 2p6 3s2 3p6 4s2 3d1. So for scandium the 1st and 2nd electron must be in 1s orbital, the 3rd and 4th in the 2s, the 5th through 10th in the 2p orbitals, etc.

In scandium example, 4s has lower energy and appears before 3d (the complexity of the d-orbital leads to its higher energy), so it is written before adding 3d to the electron configuration. But it is common to keep all the principle quantum numbers together so you may see the electron configuration written as Sc, 1s2 2s2 2p6 3s2 3p6 3d1 4s2. Writing electron configurations like this can cause difficulties in determining the element that matches an electron configuration. But if you just count the number of electrons it will equal the number of protons. This equals the atomic number which is unique for each element. For example: “Which element has the electron configuration: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d8 5s2?” Counting the electrons gives 46, which is the atomic number of palladium.

Here’s a diagram of the first several electron configurations. David’s Whizzy Periodic Table is a visual way of looking at the changing electron configuration of elements. &

Periodic Table of Elements Listing Electron Configurations.

(#e– = 1)

H 1s1 2 (#e– = 2) He 1s2 3 Li 1s2 2s1 4 Be 1s2 2s2 5 B 1s2 2s2 2p1 6 C 1s2 2s2 2p2 7 N 1s2 2s2 2p3 8 O 1s2 2s2 2p4 9 F 1s2 2s2 2p5 10 Ne 1s2 2s2 2p6 11 Na 1s2 2s2 2p6 3s1 12 Mg 1s2 2s2 2p6 3s2 13 Al 1s2 2s2 2p6 3s2 3p1 14 Si 1s2 2s2 2p6 3s2 3p2 15 P 1s2 2s2 2p6 3s2 3p3 16 S 1s2 2s2 2p6 3s2 3p4 17 Cl 1s2 2s2 2p6 3s2 3p5 18 Ar 1s2 2s2 2p6 3s2 3p6

19 K 1s2 2s2 2p6 3s2 3p6 4s1 20 Ca 1s2 2s2 2p6 3s2 3p6 4s2 31 Ga 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1 32 Ge 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p2 33 As 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3 34 Se 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p4 35 Br 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5 36 Kr 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6

21 Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1 22 Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2 23 V 1s2 2s2 2p6 3s2 3p6 4s2 3d3 24 Cr 1s2 2s2 2p6 3s2 3p6 * 4s1

  • 3d5 25
Mn 1s2 2s2 2p6 3s2 3p6 4s2 3d5 26 Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 27 Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7 28 Ni 1s2 2s2 2p6 3s2 3p6 4s2 3d8 29 Cu 1s2 2s2 2p6 3s2 3p6 * 4s1

  • 3d10 30
Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10

Note the 3d orbital follows the 4s in the lowest row, but starting with Ga (#31) it is next to the 3p orbital. It is most commonly listed with the other 3 orbitals, but sometimes it follows the 4s orbital to indicate that the 3d orbital is lower in energy than the 4s while it is being filled.

Periodic Table Exceptions To Know There is a major exception to the normal order of electron configuration at Cr (#24) and Cu (#29). It turns out that the energy the electron configuration that is half-filled, 4s1 3d5, and filled orbital, 4s1 3d10, has lower energy than the typical filling order, 4s23d4, and 4s2 3d9. This pattern is followed in the 5th row with Mo (#42) and Ag (#47).

f block Elements For completeness a couple of f-block elements are listed here. Neodynmium, Nd, which is used in very powerful magnets, has an atomic number of 60. For 60 electrons the electron configuration is: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s2 4f4. For californium, Cf, with 98 electrons the electron configuration is: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s2 4f4 Orbital Diagrams Many times it is necessary to see all the quantum numbers in an electron configuration, thus the purpose of the orbital diagram. In addition to listing the principle quantum number, n, and the subshell, l, the orbital diagram shows all the different orientations and the spin of every electron. The diagram shows the number of subshell by using boxes or lines for electrons (use three for p-orbitals, five for d-orbitals, and 7 for f-orbitals). In each box the spin of an electron is noted by using arrows. Up arrows mean ½ spin and down arrows mean –½ spin. For example, the orbital diagram for the first 18 atoms are shown below.


1s ? 2 He 1s ?? 3 Li 2s ? 1s ?? 4 Be 2s ?? 1s ?? 5 B 2p ? 2s ?? 1s ?? 6 C 2p ? ? 2s ?? 1s ?? 7 N 2p ? ? ? 2s ?? 1s ?? 8 O 2p ?? ? ? 2s ?? 1s ?? 9 F 2p ?? ?? ? 2s ?? 1s ?? 10 Ne 2p ?? ?? ?? 2s ?? 1s ?? 11 Na 3s ? 2p?? ?? ?? 2s ?? 1s ?? 12 Mg 3s ? ? 2p?? ?? ?? 2s ?? 1s ?? 13 Al 3p ? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ?? 14 Si 3p ? ? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ?? 15 P 3p ? ? ? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ?? 16 S 3p ?? ? ? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ?? 17 Cl 3p ?? ?? ? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ?? 18 Ar 3p ?? ?? ?? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ??

Rules for Filling Orbitals Aufbau Principle states that the lowest energy orbital is filled first. So electrons usually fill the lowest energy level and the simplest orbital shape first. Pauli Exclusion Principle states that no two electrons can have the same four quantum numbers. This is why each orbital only has two electrons, one spin up ( ½) and one spin down (–½). Hund’s Rule states that orbitals of the same energy, those that differ only in their orientation, are filled with electrons with the same spin before the second electron is added to any of the orbitals. This is why electrons have up spin, ?, in the orbital diagrams of B to N and of Al to P in the diagrams above.

More Examples of Orbital Diagrams Here are some orbital diagrams of elements with more electrons to help you understand the rules, electron configuration, orbital diagrams, and quantum numbers. 23 V 3d ? ? ? 4s ?? 3p ?? ?? ?? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ??

Hund’s rule requires filling orbitals that are similar with all spin the same until 1/2 filled. 47 Ag 4d ?? ?? ?? ?? ?? 5s ?__ 4p ?? ?? ?? 3d ?? ?? ?? ?? ?? 4s ?? 3p ?? ?? ?? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ??

Ag is an exception like Cu 76 Os 4d ?? ? ? ? ? 4f ?? ?? ?? ?? ?? ?? ?? 6s ?? 5p ?? ?? ?? 4d ?? ?? ?? ?? ?? 5s ?? 4p ?? ?? ?? 3d ?? ?? ?? ?? ?? 4s ?? 3p ?? ?? ?? 3s ? ? 2p ?? ?? ?? 2s ?? 1s ??

Noble Gas Abbreviation Writing out the electron configuration over and over can be tedious and shifts ones attention away from the outer electrons that are the most important electrons. So an abbreviated form of electron configurations was developed using the final column of the periodic table, the noble gases. Any element can be abbreviated except H and He, by using the noble gas with fewer electrons than the element. For example, instead of Sc : 1s2 2s2 2p6 3s2 3p6 4s2 3d1 it would be abbreviated as Ar 4s2 3d1 or Ar 3d1 4s2. For Ag the abbreviation would be: Kr?5s14d10 (see orbital diagram above), and for Os: Xe?6s2 4f14 5d6 or Xe?4f14 5d6 6s2 . Just remember that the abbreviations require that you use noble gases only and that you use a noble gas with fewer electrons. Also you can’t abbreviate a noble gas by using its symbol in brackets; that is, Ar is Ne?3s2 3p6 not Ar. Finally, you can still count the number of electrons to determine the element, you just start with the number of electrons in the noble gas. For example “What is the element with the electron configuration: Xe?6s2 4f14 5d6?” Counting the electrons 54 + 2 + 14 + 6 = 76 which is the atomic number for osmium, Os.

Similarities in Electron Configuration Equal Similar Properties. Now we can put together the first and second part of this unit. When the periodic table was being developed, chemists looked for similarities in chemical and physical properties. Any theory that describes the arrangement of electrons must be able to explain these similarities. Let’s look at the electron configurations in a periodic table format again.

Periodic Table of Elements Listing Electron Configurations.

(#e– = 1)

H 1s1 2 (#e– = 2) He 1s2 3 Li 1s2 2s1 4 Be 1s2 2s2 5 B 1s2 2s2 2p1 6 C 1s2 2s2 2p2 7 N 1s2 2s2 2p3 8 O 1s2 2s2 2p4 9 F 1s2 2s2 2p5 10 Ne 1s2 2s2 2p6 11 Na 1s2 2s2 2p6 3s1 12 Mg 1s2 2s2 2p6 3s2 13 Al 1s2 2s2 2p6 3s2 3p1 14 Si 1s2 2s2 2p6 3s2 3p2 15 P 1s2 2s2 2p6 3s2 3p3 16 S 1s2 2s2 2p6 3s2 3p4 17 Cl 1s2 2s2 2p6 3s2 3p5 18 Ar 1s2 2s2 2p6 3s2 3p6

19 K 1s2 2s2 2p6 3s2 3p6 4s1 20 Ca 1s2 2s2 2p6 3s2 3p6 4s2 31 Ga 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1 32 Ge 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p2 33 As 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3 34 Se 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p4 35 Br 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5 36 Kr 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6

21 Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1 22 Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2 23 V 1s2 2s2 2p6 3s2 3p6 4s2 3d3 24 Cr 1s2 2s2 2p6 3s2 3p6 * 4s1

  • 3d5 25
Mn 1s2 2s2 2p6 3s2 3p6 4s2 3d5 26 Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 27 Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7 28 Ni 1s2 2s2 2p6 3s2 3p6 4s2 3d8 29 Cu 1s2 2s2 2p6 3s2 3p6 * 4s1

  • 3d10 30
Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10

The first column of this periodic table has a single electron in the outer s-orbital: H 1s1, Li 2s1, Na 3s1, K 4s1 . So that similarity in the electrons in the outer energy level must be why the alkali metals are all acting the same both in their chemistry and their physical properties. Hydrogen is an exception because it is a single proton in the nucleus and a single electron which gives it wholly unique properties (although at high pressure and low temperature it can act as a metal). All the way through the periodic table we see this same pattern in each column where the outermost electrons have filled the subshells in a similar manner. For transition metals the outermost electrons are the 4s2 electrons that surround the filling 3d orbital (the 4s is in the 4th energy level and the 3d is in the lower 3rd energy level). As transition metals add more electrons and more protons their properties change more subtly than the alkali metals and alkaline earth metals, because the outermost electrons are nearly always the same (remember the exceptions of Cr and Cu).

The outermost electrons are so important that we give them a name: valence electrons. The valence electrons will be a major part of our discussion of bonding and formation of compounds. Periodic Trends We have already identified one trend on the periodic table: the change from metals to nonmetals as we move right across the groups. There are other periodic trends that show change in properties as you move across the groups or down the period. There are interactive periodic tables on the internet that allow you to pick properties and see the changes in color across the periodic table. Please visit : Dynamic Periodic Table,, and Interactive Periodic Table,

There are three major trends. Ionization energy which is the energy needed to remove an electron from an atom or an ion (an ion is an atom with electrons added or removed to give it a negative or positive charge respectively). Electronegativity which is a measure of the ability of an atom to pull electrons from another atom while the two are bonded. And atomic radius, or size, which is the ½ the distance between two atoms of an element that are bonded (less specifically, it is the radius of a single atom). The ionic radius is similar to the atomic radius, but every added electron increases the radius and every removed electron decreases the radius. Thus, all negative ions are larger than positive ions in the same period, or row, since positive ions have electrons removed.

Factors Influencing the Periodic Trends The periodic trends have their basis in the interaction of charge. The positive protons in the nucleus are attracting, or pulling, on the negative electrons around it. Also the electrons are repelling each other so that inner electrons are pushing the outer electrons away and outer electrons are repelling each other to the farthest space.

Effective nuclear charge is the positive charge pulling the outer electrons towards the nucleus. The effective nuclear charge increases as you move across a period, since each atom is gaining a new proton in the nucleus. The term “effective” is used because adding one proton does not increase attractive force by +1 charge because of the shielding effect of inner electrons in an atom. For example, the effective nuclear charge for the second period is Li = 1.28, B = 2.58, C = 3.22, N = 3.85, O = 4.49, and F = 5.13. You can see that the effective nuclear charge is not the same as the total nuclear charge: +3 for Li to +9 for F. The effective nuclear charge is distributed over the number of electrons in the atom.

Shielding or Screening is a property of inner electrons reducing the attractive force of the nucleus on the outer electrons. Shielding increases as you move down a period. Since electrons repel each other, the inner electrons keep outer electrons farther away from the nucleus. Coulombs law says that the distance charged sources apart reduced the force between the sources by 1/r, where r is 1/2 the distance between the two sources. Thus as the electrons move away from the nucleus, the force of attraction is reduced significantly. You may have experienced something similar when you move a magnet closer and closer to a metal surface all of a sudden the force increases drastically and the magnet moves against the surface it is attracted to in a little jump. So if the outer electrons are pushed away then the attractiveness is reduced in a measurable amount.

Ionization Energy a Proof of the Quantum Model of Electron Arrangement Ionization energy is the energy needed to pull an electron away from an atom. Ionization energy is highest for helium, He and lowest for francium, Fr.

Ionization Energy highlighting Alkali Metals and Noble Gases

The figure above shows the trend in ionization. For families like alkali metals and noble gases, the ionization energy decreases by moving down a group. This is because each time an element has a new energy level, or shell added, the electrons are farther away from the nucleus. These outer electrons are shielded by the inner electrons. Ionization energy also increases moving across a group from left to right. This is because the number of protons and the positive charge in the nucleus increases across the period. So the electrons are attracted more strongly to the nucleus (and there is no additional shielding because as you move across a period you don’t change inner electrons) .

A more interesting observation of the ionization energy graph is how it mirrors the quantum theory of electron arrangement. Noble gases all have the highest ionization energy in their period and the very next element, which is an alkali metal, has a lower ionization energy about 2 to 4 times smaller. This observation matches our model of a new larger electron shell being started for the alkali metals.

Furthermore, the ionization energy from Li to Be increases but for B decreases a bit. This confirms the change from the s-orbital to p-orbital and the higher energy of a p-orbital that would make the electron easier to remove.

Then the ionization energy increases from B to N as more protons in the nucleus increase the attractive force and make the electrons more difficult to remove. But for N to O there is another decrease in ionization energy. This matches our theory that the extra electron in oxygen (compared to nitrogen) is placed in an orbital containing another electron and the repelling force between the two electrons in the same orbital reduces the ionization energy.

Finally, we can see the long flattened area of the graph to the right of K, potassium, which is the d-orbital electrons. Placed beneath the 4s electrons, the 3d electrons provide shielding that cancel out the added proton to the nucleus, and the valence electrons in 4s have the same ionization energy across the transition metal.


Electronegativity, which is a measure of the attractive force of an atom’s nucleus for the electrons that are shared in a bond with another atom, follows the same trend as ionization energy. Electronegativity is highest in the upper right with F, fluorine and lowest in the lower left with Fr, francium. However, since noble gases are unreactive and they don’t bond with another atom usually, there is no electronegativity for He, Ne, and Ar.

Periodic Trends modified from

Atomic Radius or Atomic Size The atomic radius of an atom will be smaller when there is a greater positive charge in the nucleus. There are fewer electron-electron repulsive forces (shielding), and the number of energy levels, or shells, is greater. Thus, the trend for atomic radius is opposite to the trend for ionization energy; the atomic radius increases as you go down a group (adding energy levels) and increases from right to left across a period.

Ionic Radius or size of ion The ions of elements form when a chemical reaction causes an electron to be added or removed from a neutral element (neutral is when the number of electrons and protons are equal). Ions that gain electrons have an increased radius, since electron-electron repulsions increase and the attractive force of the nucleus is distributed over a larger number of negative charges. Ions that lose electrons on the other hand will shrink because there are fewer repulsive forces and the positive, attractive force acts on fewer negative electrons. So while the trend in ionic size is the same as for atomic size (increasing size moving down and to the right), the ions that have gained electrons (called anions), are always larger than the ions that have lost electrons (called cations).

Summary Unit 1. Pure substances can be classified as elements and compounds. Elements are made of atoms and our understanding of the nature of the atom has evolved over two centuries. The atom is composed of a nucleus that is extremely small compared to the outer boundary of the atom, but contains nearly all the mass of the atom (so it is super dense). The nucleus contains protons and neutrons. The number of protons is unique to each element and is the same as the atomic number on the periodic table. The number of neutrons changes within atoms of the same element to make different isotopes of the element. The mass number is the number of protons and neutrons in each isotope, while the periodic table lists the average atomic mass of the element (measured in amu, atomic mass units). The average atomic mass is determined by using the percent natural abundance of each isotope of the element and the mass of each isotope in a calculation called a weighted average. Electrons surround the nucleus and form the boundary of the atom called the electron cloud. The number of electrons equals the number of protons for a neutral atom, because the negative charges of the electrons match the positive charge of the protons in the nucleus. Changing the number of electrons forms an ion of the atom. Gaining electrons increases the negative cha

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