simple schooling the periodic table by j. anne huss - ©...

Simple Schooling The Periodic Table By J. Anne Huss - © 2011

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Elements have always been a part of the human experience. Even in ancient times gold and silver were revered for their physical properties such as high luster and ease of malleability, allowing ancient people to fashion it into beautiful jewelry, coins, and other decorative objects. These same physical properties are what scientists use to understand chemical reactions in the natural world.

What these ancient people did not know is that the elements were all around them and that everything is made up of the elements. In 330 B.C. Aristotle proposed that the elements belonged to one of the four roots, but Plato later renamed these groups elements. The elemental groups were known as earth, air, fire, and water. The earth group consisted of things which were first and foremost dry, but secondarily cold. The air group contained

elements which were thought of as primarily wet, but secondarily hot. The fire group contained things which were primarily hot, but secondarily dry, and the water group contained things which were primarily

cold and secondarily wet. If you think about this organization it doesn’t really make a lot of sense but that was all “scientists” had for more than a thousand years. Finally, in 1817 Johann Wolfgang Döbereiner came up with another way to do things. He hypothesized that some of the known elements would form triads with related properties. Most of this was based on the average of the three atomic weights of each element.


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Döbereiner discovered that the atomic weight of the first and third element was almost the exactly the same as the third element. He found a pattern. These triads were organized like this:

As time went on more and more elements were discovered and by 1869 there were 63. The more elements there were, the easier it is to see a pattern, and when a pattern is found the natural thing to do is to try and make sense of it. This is exactly what chemists of the 19th century did. There were a few early attempts at elemental classification, but it would be Dmitri Mendeleev who would essentially crack the code and give us the periodic table of the elements as we know it now.

Mendeleev was born to a family with a moderate amount of money in Russia in 1834 but ended up poor and living in St. Petersburg due to his father’s death and a fire which destroyed his mother’s factory. He studied science at several universities (in St. Petersburg and Germany) and became a professor of chemistry in 1863 at St. Petersburg State University. He spent the next several years pondering the periodicity of the elements. Periodicity means things that occur at regular intervals. He was not the first scientist to think about this – in fact several others, including Alexandre-Emile Béguyer de


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Chancourtois in 1862 and John Newlands in 1865, had already proposed that the elements could be arranged by atomic weight and in groups of eight. However, neither of these two could connect all the dots that Mendeleev eventually would.

Mendeleev began by arranging the elements in a table ordered by atomic weight, but went on to explain how and why he was doing so in a paper published in 1869, both in Russia and Germany. This

paper identified eight observations about how elements could be classified. The first observation states: The elements, if arranged according to their atomic weights, exhibit an evident periodicity of properties. Remember that periodicity means that things will occur over and over again in regular intervals. So Mendeleev's first observation states that certain characteristics of the elements repeat over and over again. The second observation states: Elements which have similar chemical properties have atomic weights which are either of nearly the same value (e.g., platinum, iridium, osmium) or which increase regularly (e.g., potassium, rubidium, cesium). The third observation states: The arrangement of the elements, or of groups of elements in the order of their atomic weights corresponds to their valencies as well as to their distinctive chemical properties.


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In this observation Mendeleev proposes that the atomic weights are correlated with the number of valence electrons. Of course, electrons were not known in that time, but chemists were beginning to figure things out. Mendeleev’s table would propel discovery forward in the areas of chemistry and physics. The fourth and fifth observations state: The elements which are the most widely diffused have small atomic weights. The value of the atomic weight determines the character of the element.

Observations six through eight state: We must expect the discovery of many yet unknown elements. The atomic weight of an element may sometimes be amended by a knowledge of those of the elements around it. Certain characteristic properties of the elements can be predicted from their atomic weights. These 8 ideas became known as Mendeleev’s eight periodic principles and they allowed the known elements to be arranged in a table according to their atomic weights and similar physical properties. Since there were only 63 known elements on that original periodic table there were gaps in the table where larger than expected atomic weight jumps occurred. Mendeleev hypothesized that these elements existed, but were as of yet – unknown to man. In fact, he predicted the existence of gallium, scandium, and germanium – even though no one knew of these elements at the time of the prediction. Years later (but still during Mendeleev’s lifetime) these elements were


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discovered and the credibility of his periodic table of the elements was raised to accepted status in the scientific community.

Even though the table answered many things about elements at the time, there were still some questions. For instance, why are we able to arrange some into the table according to increasing mass, but some others did not quite fit the pattern? In addition, why did the elements exhibit periodicity? Recall that periodicity means “to

occur at regular intervals” – and in fact, the properties of the elements also occurred at regular intervals. 40 years after Mendeleev’s periodic table a scientist named Henry Moseley discovered another elemental pattern. Moseley was


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examining the x-ray spectra of various elements when he observed that the elements of the periodic table lined up consistently when arranged according to the number of positive charges in the nucleus. The positive charge found inside the nucleus of the atom is of course, the proton. And that is why today’s periodic table is arranged according to number of protons in the nucleus of the atom and not strictly by atomic weight. Until Moseley's work, the "atomic number" was just an element's place in the periodic table, and was not known to be associated with any measureable physical quantity. Moseley was able to come up with a mathematical formula which connected the frequencies of x-rays emitted from the elements to the square of a number which was close to that element’s atomic number. This became known as Moseley’s Law. The modern periodic table is divided up into vertical groups and horizontal periods as well as according to type of element. For example – the transition metals are in the middle, the noble gases are on the far right, and the halogens are just left of the noble gases. Scientists didn’t arbitrarily organize the table this way – remember that they are arranged according to atomic weight and number of protons in the nucleus (atomic number). But in addition to these reasons, the elements are categorized this way because they share physical properties.


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The periodic table can be divided into three major types of elements:

• Metals • Non-metals • Semi-metals

Metals are elements which typically have fewer electrons available for chemical reactions. They also tend to want to lose electrons, causing them to become positive. In addition, the metals can be good conductors of electricity and heat. Most of the elements on the table are metals and examples include gold, silver, platinum, iron, cobalt, manganese, magnesium, and lead. Non-metals are elements which have more electrons available for chemical reactions and also to gain electrons – which gives them an overall negative charge. They do not conduct electricity or heat very well. Non-metal elements include helium, hydrogen, carbon, oxygen, chlorine, sulfur, and iodine. The semi-metals are also called metalloids and will often have a few characteristics of both metals and non-metals. A few examples of semi-metals include boron, arsenic, and silicon.


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Notice the groups labeled 1-18 running along the top of the table. These groups represent the elements arranged in vertical columns according to chemical properties. That means that all the elements in group 14 share similar properties, all the elements in group 18 share similar properties, and so on. Also notice periods 1-7 labeled on the left side of the table. Elements of the same period have the same number of electron shells; with each

group across a period, the elements have one more proton and electron and become less metallic. For practical purposes an electron shell can be thought of as an “orbit” or path followed by electrons as they move around the atom’s nucleus. The number of shells depends on the element’s place in the periodic table. The shell closest to the nucleus is called the 1 shell, followed by the 2 shell, 3 shell and so on.

Each shell can contain only a fixed number of electrons and scientists use this fact to predict behavior in chemical reactions. Knowing the period of an element is useful because all elements in the same period will have their valence electrons in the same energy level. The valance electrons are the outermost electrons in an element. These are the


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electrons which are capable of creating chemical reactions, depending on how many are available. Not all valence electrons are available. It turns out that 8 is a magic number in chemistry because 8 valence electrons represent the maximum number of electrons capable of being in the outer most energy level (shell). When the outer most energy level, or shell, has all 8 electrons it is called a “full shell”. If the outer most shell has less than 8 electrons it is called an “open shell”. We can tell how many valence electrons are in some of the elements just by looking at group numbers. Study the diagram below and notice that the group numbers have been replaced with the number of valence electrons:


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If you look at the first periodic table where group numbers were introduced and compare it to the valence table you will notice that group numbers 3-12 have been removed (the transition metals) and the tens place has been removed from the group numbers 13-18. This leaves us with 8 groups and each of these numbers represents the number of valence electrons in those elements in that group. For example carbon (C) has 4 valence electrons, boron (B) has 3 valence electrons, and Hydrogen (H) has 1 valence electron. The only element which does not follow this rule is helium, and that is because it only has two electrons to begin with, thus it has a full shell at 2. The magic number of valence electrons is 8, this means that when an element has 8 valence electrons it is happy and does not want to create a chemical reaction or

find any more electrons to fill its shell. If, however, an element does not have 8 valence electrons it seeks other elements out to fill up the missing spots. For example, oxygen has 6 valence electrons but wants 8 – so what happens is that the free valence electrons will look for another atom to bond with to fill up the open shell.


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The noble gases (helium, neon, argon, krypton, xenon, and radon) are all in the new group 8, so this means that they have a full shell and are not “reactive”. They don’t combine with other elements very easily. In fact it is quite difficult to get these elements to do much of anything. They are called “stable”.

Hydrogen, on the other hand, has only 1 valence electron – making it VERY reactive! It wants to combine with just about everything. This makes it a very unstable element in its pure form. Sodium, is in the same GROUP as hydrogen so thus has only 1 valence electron, also wants to combine with just about everything. In fact, pure elemental sodium is so reactive it cannot even be exposed to air because of water content in the air (humidity) is enough to cause it to explode! Notice that the elements hydrogen and sodium are both alkali metals and belong to the same group. This means that other elements in that group will also have similar reactive properties. Potassium (K) is also very explosive when exposed in its pure form to air. Elements which behave in this fashion must be stored under special conditions – in the case of potassium it

must be stored in kerosene or oil to prevent any possible chemical reaction with water. So far you have learned that the periodic table is organized according to: Number of protons

• Atomic weight • Groups • Periods

But that’s not where it ends.


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In addition to those above there are also “trends” in the periodic table which relate to:

• Atomic radii • Ionization energy • Electron affinity • Ionic radii • Electronegativity

Study the table of trends carefully before moving on.

Atomic radius describes the distance from the nucleus to the boundary of the electron cloud. Recall that the electrons are not really orbiting the nucleus like planets orbit the sun, but instead occupy an area some distance from the nucleus in a cloud fashion. This is not a hard and fast measurement because the electron cloud is not a well defined area, but is instead a place where electrons could be found. The trend for atomic radius states that as you move down a group the atomic radius will increase because the number of energy levels in an atom also increases.


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Each energy level, which is like the “orbits” portrayed in older models of the atom, will be farther and farther away from the nucleus as the number of protons and electrons increase. The second part of the atomic radius trend is that as you move across a period, the atomic radius will decrease. This is because no new energy levels are added to the elements as you move across a period and there is a stronger force of attraction towards the nucleus, making the atomic radius smaller. The ionization energy is the amount of energy needed to remove an electron off of an atom. The first electron to be removed is the easiest, but each subsequent electron requires more and more energy as the nucleus becomes more and more positive. A low ionization energy means that it takes a small amount of energy to remove electrons from the atom.


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Electron affinity is the energy given off when an atom in the gas phase gains extra electrons to form a negatively charged ion.

The ionic radii is similar to the atomic radii, with the exception that ions have charges on them, meaning they either have more or fewer electrons than a neutral atom. An atom with extra electrons is called an anion and has a net negative charge. An atom that is missing some electrons is called a cation and has a net positive charge.

Anions are larger than cations or neutral atoms because the negative charge on the electrons has a repulsion force which pushes them apart, this causes the electron cloud to be larger. Additionally, anions have more electrons than protons so the protons cannot pull the electrons towards the nucleus like they do on neutral atoms or cations.


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The trend for atomic radii is the same as it was for ionization energy – it gets larger as you move across periods and as you move up groups.

Electronegativity is the measure of the ability of an atom to attract pairs of electrons to create a chemical bond. The trend for electronegativity is that as you move from right to left the electronegativity decreases. Moving moving down a group will also cause the electronegativity to decrease.


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The periodic table is a tool for you to use in your study of chemistry, biology, and physics. It tells you many things and learning how it is organized will make your future study of science easier and more productive. Understanding it now during your study of biology will greatly enhance your ability to grasp the more complicated topics in chemistry. The periodic table, when properly understood, is a great cheat sheet to understanding the world around you!


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Student Activities Exercise One Write the letter of the correct match next to each problem.

1. Element a. Proposed the law of octaves to describe the

periodicity of the chemical elements.

2.

Johann Wolfgang Döbereiner

b. The repetition of similar properties in chemical elements, as indicated by their positioning in the periodic table.

3.

Triads

c. A mathematical formula that connects the frequencies of x-rays emitted from the element to the square of a number which is close to that element’s atomic number.

4. Dmitri Mendeleev

d. A German chemist who is arranged the chemical elements into triads.

5. Periodicity e. Formulated a periodic table similar to the

one we use today.

6.

John Newlands

f. A positively charged particles which resides inside the nucleus of an atom and the number of which convey the elements identity.

7. Henry Moseley g. A group of three.

8.

Proton

h. Correlated the positive charges in the nucleus to the atomic number on the periodic table.

9. Moseley’s Law i. A substance composed of atoms having an

identical number of protons in each nucleus.

10.

Period

j. A sequence of elements arranged in order of increasing atomic number and forming one of the horizontal rows in the periodic table.


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Exercise Two Write the letter of the correct match next to each problem.

1. Group a. Elements which exhibit properties of both

metals and non-metals.

2. Noble gases b. Elements which typically have fewer electrons

available for chemical reactions.

3. Halogens c. Negatively charges particles which “orbit” the

nucleus of an atom.

4. Metals d. An orbit followed by electrons around an atom

nucleus.

5. Electrons e. Reactive elements with 7 valence electrons.

6.

Conductor

f. The electrons of an atom that can participate in the formation of chemical bonds with other atoms.

7. Non-metals g. An element which allows for free flow of

electrons and heat.

8. Metalloids h. Includes helium, neon, argon, krypton, xenon,

and radon.

9. Electron

shell

i. A vertical column of elements in the periodic table that all have similar electronic structures, properties, and valencies.

10.

Valence electrons

j. Elements which have more electrons available for chemical reactions and which tend to gain electrons giving them an overall negative charge.


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Exercise Three Write the letter of the correct match next to each problem.

1. Full shell a. The distance from the nucleus to the edge

of the electron cloud.

2. Open shell b. A shell which has the maximum number of

electrons.

3. Trend c. A shell which is not filled and is capable of

chemical bonding.

4.

Atomic radii

d. The energy given off when an atom in the gas phase gains extra electrons to form a negatively charged ion.

5. Ionization energy

e. A negatively charged atom due to extra electrons.

6. Electron affinity f. Can be larger or smaller than atomic radii.

7. Ionic radii g. The general direction in which something

tends to move.

8.

Anion

h. The tendency for an atom to attract a pair of electrons that it shares with another atom.

9. Cation i. The amount of energy needed to remove

an electron off of an atom.

10. Electronegativity j. A positively charged atom due to missing

electrons.


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Exercise Four Label the major sections of the periodic table using the word bank below.

Word Bank

Metalloids Non-metals

Metals


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Exercise Five Label ALL of the sections of the periodic table using the word bank below.

Word Bank Lanthinides & Actinides

Non-Metals Metals

Alkali Metals Alkaline Earth Metals

Metalloids Noble Gases Synthetics

Transitions Metals


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Exercise Six Fill in the blanks by using your periodic table to answer the questions. 1. What is the name of the third element from the top of group eleven. 2. What is the name of the fourth element from the right in period four. 3. Lead is the fifth element from the top of group _______. 4. Potassium is the _____ element in period ______. 5. This period contains carbon, oxygen, and nitrogen.

You can also use this printable periodic table if you have trouble seeing the small print.

http://profmokeur.ca/chemistry/printable_periodic_tablecol.pdf�


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Exercise Seven Final Test Choose the best answer to each question.

1. Johann Wolfgang Döbereiner proposed: a. The Law of Octaves b. Elements form triads c. The proton = atomic number

2. Dmitri Mendeleev proposed:

a. The Law of Octaves b. The element gallium c. The proton = atomic number

3. Henry Moseley proposed:

a. The element gallium b. Elements form triads c. The proton = atomic number

4. Periods are arranged according to:

a. Increasing atomic number b. Valence electrons c. Open shells

5. Which element is in period four group one?

a. Potassium b. K c. Both of these

6. Which element is in period one group one?

a. Helium b. Hydrogen c. None of these

7. Which element is in group fourteen period one?

a. Carbon b. Nitrogen c. Neither of these


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8. Which element is in group eighteen period two? a. Hydrogen b. Neon c. Argon

9. Groups are arranged according to:

a. Valencies b. Electrons available for bonding c. Both of these

10. In all cases the ____ equals the atomic number.

a. Proton b. Neutron c. Electron

11. The ______ is what enables atoms to react with other atoms.

a. Neutrons b. Protons c. Electrons

12. The pure element gold has atoms that have how many protons?

a. 79 b. 196.9 c. Both of these

13. Atomic radii measures the distance between the neutron and the ______.

a. Edge of the atom b. Edge of electron cloud c. The nearest electron

14. Electronegativity measures the ability of an atom to attract ______.

a. Pairs of electrons b. Pairs of protons c. Pairs of neutrons


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15. An atom with an overall net positive charge is called a/an ________. a. Cation b. Anion c. Ionic radii


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Parent Solutions Exercise One Write the letter of the correct match next to each problem.

1. i Element a. Proposed the law of octaves to describe the

periodicity of the chemical elements.

2. d

Johann Wolfgang Döbereiner

b. The repetition of similar properties in chemical elements, as indicated by their positioning in the periodic table.

3.

g

Triads

c. A mathematical formula that connects the frequencies of x-rays emitted from the element to the square of a number which is close to that element’s atomic number.

4. e Dmitri Mendeleev

d. A German chemist who is arranged the chemical elements into triads.

5. b Periodicity e. Formulated a periodic table similar to the

one we use today.

6. a

John Newlands

f. A positively charged particles which resides inside the nucleus of an atom and the number of which convey the elements identity.

7. h Henry Moseley g. A group of three.

8. f

Proton

h. Correlated the positive charges in the nucleus to the atomic number on the periodic table.

9. c Moseley’s Law i. A substance composed of atoms having an

identical number of protons in each nucleus.

10. j

Period

j. A sequence of elements arranged in order of increasing atomic number and forming one of the horizontal rows in the periodic table.


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Exercise Two Write the letter of the correct match next to each problem.

1. i Group a. Elements which exhibit properties of both

metals and non-metals.

2. h Noble gases b. Elements which typically have fewer electrons

available for chemical reactions.

3. e Halogens c. Negatively charges particles which “orbit” the

nucleus of an atom.

4. b Metals d. An orbit followed by electrons around an atom

nucleus.

5. c Electrons e. Reactive elements with 7 valence electrons.

6. g

Conductor

f. The electrons of an atom that can participate in the formation of chemical bonds with other atoms.

7. j Non-metals g. An element which allows for free flow of

electrons and heat.

8. a Metalloids h. Includes helium, neon, argon, krypton, xenon,

and radon.

9. d Electron

shell

i. A vertical column of elements in the periodic table that all have similar electronic structures, properties, and valencies.

10.

f Valence electrons

j. Elements which have more electrons available for chemical reactions and which tend to gain electrons giving them an overall negative charge.


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Exercise Three Write the letter of the correct match next to each problem.

1. b Full shell a. The distance from the nucleus to the edge

of the electron cloud.

2. c Open shell b. A shell which has the maximum number of

electrons.

3. g Trend c. A shell which is not filled and is capable of

chemical bonding.

4. a

Atomic radii

d. The energy given off when an atom in the gas phase gains extra electrons to form a negatively charged ion.

5. i Ionization energy

e. A negatively charged atom due to extra electrons.

6. d Electron affinity f. Can be larger or smaller than atomic radii.

7. f Ionic radii g. The general direction in which something

tends to move.

8. e

Anion

h. The tendency for an atom to attract a pair of electrons that it shares with another atom.

9. j Cation i. The amount of energy needed to remove

an electron off of an atom.

10. h Electronegativity j. A positively charged atom due to missing

electrons.


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Exercise Four

Exercise Five


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Exercise Six Fill in the blanks by using your periodic table to answer the questions. 1. What is the name of the third element from the top of group eleven. Answer: Gold 2. What is the name of the fourth element from the right in period four. Answer: Arsenic 3. Lead is the fifth element from the top of group _______. Answer: Fourteen 4. Potassium is the _____ element in period ______. Answer: first, four 5. This period contains carbon, oxygen, and nitrogen. Answer: two


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Exercise Seven Final Test Choose the best answer to each question.

1. Johann Wolfgang Döbereiner proposed: a. The Law of Octaves b. Elements form triads c. The proton = atomic number

2. Dmitri Mendeleev proposed:

a. The Law of Octaves b. The element gallium c. The proton = atomic number

3. Henry Moseley proposed:

a. The element gallium b. Elements form triads c. The proton = atomic number

4. Periods are arranged according to:

a. Increasing atomic number b. Valence electrons c. Open shells

5. Which element is in period four group one?

a. Potassium b. K c. Both of these

6. Which element is in period one group one?

a. Helium b. Hydrogen c. None of these

7. Which element is in group fourteen period one?

a. Carbon b. Nitrogen c. Neither of these


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8. Which element is in group eighteen period two? a. Hydrogen b. Neon c. Argon

9. Groups are arranged according to:

a. Valencies b. Electrons available for bonding c. Both of these

10. In all cases the ____ equals the atomic number.

a. Proton b. Neutron c. Electron

11. The ______ is what enables atoms to react with other atoms.

a. Neutrons b. Protons c. Electrons

12. The pure element gold has atoms that have how many protons?

a. 79 b. 196.9 c. Both of these

13. Atomic radii measures the distance between the neutron and the ______.

a. Edge of the atom b. Edge of electron cloud c. The nearest electron

14. Electronegativity measures the ability of an atom to attract ______.

a. Pairs of electrons b. Pairs of protons c. Pairs of neutrons


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15. An atom with an overall net positive charge is called a/an ________. a. Cation b. Anion c. Ionic radii

simple schooling the periodic table by j. anne huss - ©...

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