Periodic Trends - MCAT Physical
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Which of the following would have the greatest atomic radius?
Which of the following would have the greatest atomic radius?
Atomic radius increases down each group of the periodic table and toward the left of each period. Since the elements listed are all in the same group, iodine would have the greatest atomic radius because it farther down the period compared to the others.
Atomic radius increases down each group of the periodic table and toward the left of each period. Since the elements listed are all in the same group, iodine would have the greatest atomic radius because it farther down the period compared to the others.
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Electronegativity is an important concept in physical chemistry, and often used to help quantify the dipole moment of polar compounds. Polar compounds are different from those compounds that are purely nonpolar or purely ionic. An example can be seen by contrasting sodium chloride, NaCl, with an organic molecule, R-C-OH. The former is purely ionic, and the latter is polar covalent.
When comparing more than one polar covalent molecule, we use the dipole moment value to help us determine relative strength of polarity. Dipole moment, however, is dependent on the electronegativity of the atoms making up the bond. Electronegativity is a property inherent to the atom in question, whereas dipole moment is a property of the bond between them.
For example, oxygen has an electronegativity of 3.44, and hydrogen of 2.20. In other words, oxygen more strongly attracts electrons when in a bond with hydrogen. This leads to the O-H bond having a dipole moment.
When all the dipole moments of polar bonds in a molecule are summed, the molecular dipole moment results, as per the following equation.
Dipole moment = charge * separation distance
A scientist is studying flourine. Which of the following is a possible electronegativity value for flourine?
Electronegativity is an important concept in physical chemistry, and often used to help quantify the dipole moment of polar compounds. Polar compounds are different from those compounds that are purely nonpolar or purely ionic. An example can be seen by contrasting sodium chloride, NaCl, with an organic molecule, R-C-OH. The former is purely ionic, and the latter is polar covalent.
When comparing more than one polar covalent molecule, we use the dipole moment value to help us determine relative strength of polarity. Dipole moment, however, is dependent on the electronegativity of the atoms making up the bond. Electronegativity is a property inherent to the atom in question, whereas dipole moment is a property of the bond between them.
For example, oxygen has an electronegativity of 3.44, and hydrogen of 2.20. In other words, oxygen more strongly attracts electrons when in a bond with hydrogen. This leads to the O-H bond having a dipole moment.
When all the dipole moments of polar bonds in a molecule are summed, the molecular dipole moment results, as per the following equation.
Dipole moment = charge * separation distance
A scientist is studying flourine. Which of the following is a possible electronegativity value for flourine?
Flourine must have an electronegativity value higher than oxygen. Remember your periodic trends: electronegativity increases as we move up and to the right on the periodic table. We are told in the passage that the electronegativity of oxygen is 3.44, therefore, our answer must be 3.9.
Flourine must have an electronegativity value higher than oxygen. Remember your periodic trends: electronegativity increases as we move up and to the right on the periodic table. We are told in the passage that the electronegativity of oxygen is 3.44, therefore, our answer must be 3.9.
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Electronegativity is an important concept in physical chemistry, and often used to help quantify the dipole moment of polar compounds. Polar compounds are different from those compounds that are purely nonpolar or purely ionic. An example can be seen by contrasting sodium chloride, NaCl, with an organic molecule, R-C-OH. The former is purely ionic, and the latter is polar covalent.
When comparing more than one polar covalent molecule, we use the dipole moment value to help us determine relative strength of polarity. Dipole moment, however, is dependent on the electronegativity of the atoms making up the bond. Electronegativity is a property inherent to the atom in question, whereas dipole moment is a property of the bond between them.
For example, oxygen has an electronegativity of 3.44, and hydrogen of 2.20. In other words, oxygen more strongly attracts electrons when in a bond with hydrogen. This leads to the O-H bond having a dipole moment.
When all the dipole moments of polar bonds in a molecule are summed, the molecular dipole moment results, as per the following equation.
Dipole moment = charge * separation distance
Electronegativity is associated with another function, electron affinity. What is true of electron affinity?
Electronegativity is an important concept in physical chemistry, and often used to help quantify the dipole moment of polar compounds. Polar compounds are different from those compounds that are purely nonpolar or purely ionic. An example can be seen by contrasting sodium chloride, NaCl, with an organic molecule, R-C-OH. The former is purely ionic, and the latter is polar covalent.
When comparing more than one polar covalent molecule, we use the dipole moment value to help us determine relative strength of polarity. Dipole moment, however, is dependent on the electronegativity of the atoms making up the bond. Electronegativity is a property inherent to the atom in question, whereas dipole moment is a property of the bond between them.
For example, oxygen has an electronegativity of 3.44, and hydrogen of 2.20. In other words, oxygen more strongly attracts electrons when in a bond with hydrogen. This leads to the O-H bond having a dipole moment.
When all the dipole moments of polar bonds in a molecule are summed, the molecular dipole moment results, as per the following equation.
Dipole moment = charge * separation distance
Electronegativity is associated with another function, electron affinity. What is true of electron affinity?
Chlorine has a great thermodynamic desire to capture an electron, thus taking on the electronic structure of a stable noble gas. This causes chlorine to release energy when it captures an electron as it becomes more stable.
Sodium, on the other hand, would prefer to lose an electron and gain the configuration of a noble gas. Adding an electron would however award some stability to sodium, due to the complete s orbital that this would ensue.
Second electron affinity is usually encountered for such elements as oxygen and sulfur, which form anions with the addition of two electrons. The first electron affinity gives you O- or S-, and so it takes significant energy to add another electron to an already negative ion.
Chlorine has a great thermodynamic desire to capture an electron, thus taking on the electronic structure of a stable noble gas. This causes chlorine to release energy when it captures an electron as it becomes more stable.
Sodium, on the other hand, would prefer to lose an electron and gain the configuration of a noble gas. Adding an electron would however award some stability to sodium, due to the complete s orbital that this would ensue.
Second electron affinity is usually encountered for such elements as oxygen and sulfur, which form anions with the addition of two electrons. The first electron affinity gives you O- or S-, and so it takes significant energy to add another electron to an already negative ion.
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Electronegativity is an important concept in physical chemistry, and often used to help quantify the dipole moment of polar compounds. Polar compounds are different from those compounds that are purely nonpolar or purely ionic. An example can be seen by contrasting sodium chloride, NaCl, with an organic molecule, R-C-OH. The former is purely ionic, and the latter is polar covalent.
When comparing more than one polar covalent molecule, we use the dipole moment value to help us determine relative strength of polarity. Dipole moment, however, is dependent on the electronegativity of the atoms making up the bond. Electronegativity is a property inherent to the atom in question, whereas dipole moment is a property of the bond between them.
For example, oxygen has an electronegativity of 3.44, and hydrogen of 2.20. In other words, oxygen more strongly attracts electrons when in a bond with hydrogen. This leads to the O-H bond having a dipole moment.
When all the dipole moments of polar bonds in a molecule are summed, the molecular dipole moment results, as per the following equation.
Dipole moment = charge * separation distance
Electronegativity is closely associated with the principle of ionization energy. Which of the following defines ionization energy?
Electronegativity is an important concept in physical chemistry, and often used to help quantify the dipole moment of polar compounds. Polar compounds are different from those compounds that are purely nonpolar or purely ionic. An example can be seen by contrasting sodium chloride, NaCl, with an organic molecule, R-C-OH. The former is purely ionic, and the latter is polar covalent.
When comparing more than one polar covalent molecule, we use the dipole moment value to help us determine relative strength of polarity. Dipole moment, however, is dependent on the electronegativity of the atoms making up the bond. Electronegativity is a property inherent to the atom in question, whereas dipole moment is a property of the bond between them.
For example, oxygen has an electronegativity of 3.44, and hydrogen of 2.20. In other words, oxygen more strongly attracts electrons when in a bond with hydrogen. This leads to the O-H bond having a dipole moment.
When all the dipole moments of polar bonds in a molecule are summed, the molecular dipole moment results, as per the following equation.
Dipole moment = charge * separation distance
Electronegativity is closely associated with the principle of ionization energy. Which of the following defines ionization energy?
Even though adding an electron would generate an ion (anion), ionization energy is defined as the energy needed to remove an electron.
Even though adding an electron would generate an ion (anion), ionization energy is defined as the energy needed to remove an electron.
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Electronegativity is an important concept in physical chemistry, and often used to help quantify the dipole moment of polar compounds. Polar compounds are different from those compounds that are purely nonpolar or purely ionic. An example can be seen by contrasting sodium chloride, NaCl, with an organic molecule, R-C-OH. The former is purely ionic, and the latter is polar covalent.
When comparing more than one polar covalent molecule, we use the dipole moment value to help us determine relative strength of polarity. Dipole moment, however, is dependent on the electronegativity of the atoms making up the bond. Electronegativity is a property inherent to the atom in question, whereas dipole moment is a property of the bond between them.
For example, oxygen has an electronegativity of 3.44, and hydrogen of 2.20. In other words, oxygen more strongly attracts electrons when in a bond with hydrogen. This leads to the O-H bond having a dipole moment.
When all the dipole moments of polar bonds in a molecule are summed, the molecular dipole moment results, as per the following equation.
Dipole moment = charge * separation distance
Electronegativity is closely associated with the principle of ionization energy. Which of the following is true of second, third, and successive ionization energies?
Electronegativity is an important concept in physical chemistry, and often used to help quantify the dipole moment of polar compounds. Polar compounds are different from those compounds that are purely nonpolar or purely ionic. An example can be seen by contrasting sodium chloride, NaCl, with an organic molecule, R-C-OH. The former is purely ionic, and the latter is polar covalent.
When comparing more than one polar covalent molecule, we use the dipole moment value to help us determine relative strength of polarity. Dipole moment, however, is dependent on the electronegativity of the atoms making up the bond. Electronegativity is a property inherent to the atom in question, whereas dipole moment is a property of the bond between them.
For example, oxygen has an electronegativity of 3.44, and hydrogen of 2.20. In other words, oxygen more strongly attracts electrons when in a bond with hydrogen. This leads to the O-H bond having a dipole moment.
When all the dipole moments of polar bonds in a molecule are summed, the molecular dipole moment results, as per the following equation.
Dipole moment = charge * separation distance
Electronegativity is closely associated with the principle of ionization energy. Which of the following is true of second, third, and successive ionization energies?
As you ionize atoms, you generate charged ionic species. These charges will resist further ionization (cations will more strongly attract the electrons you are trying to pull away). Noble gas configurations are particularly stable, so you would expect a large increase in needed energy to ionize away from this state.
As you ionize atoms, you generate charged ionic species. These charges will resist further ionization (cations will more strongly attract the electrons you are trying to pull away). Noble gas configurations are particularly stable, so you would expect a large increase in needed energy to ionize away from this state.
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Which of the following elements has the greatest atomic radius?
Which of the following elements has the greatest atomic radius?
Atomic radius can be determined using the periodic trends. Atomic radius increases to the left of a period and down a group of the periodic table. Electronegativity, in contrast, increases to the right of a period and up a group of the periodic table. Relating the two, we can see that the greater the atomic radius, the weaker its electronegativity because the electrons are farther away from the nucleus and are unable to feel the attractive force of the protons in the nucleus.
Atomic radius can be determined using the periodic trends. Atomic radius increases to the left of a period and down a group of the periodic table. Electronegativity, in contrast, increases to the right of a period and up a group of the periodic table. Relating the two, we can see that the greater the atomic radius, the weaker its electronegativity because the electrons are farther away from the nucleus and are unable to feel the attractive force of the protons in the nucleus.
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Using your knowledge of periodic trends, which statement is incorrect?
Using your knowledge of periodic trends, which statement is incorrect?
The second ionization energy of a neutral atom is always greater than the first.
Recall that ionization energy refers to the energy required to remove an electron from an atom. The first ionization energy will result in a cation. The reduced number of electrons in a cation allows them to be pulled closer to the nucleus by the positively charged protons, effectively decreasing the atomic radius and increasing the attractive force between the electrons and the nucleus. The energy required to remove a second electron must overcome this increased affinity, and will thus be greater than the first ionization energy.
The rest of the answers can be examined with a few simple trends kept in mind. Atomic radius increases to the left of a period and down a group. If two ions have the same electron configuration (like the chlorine anion and potassium cation), then the anion will have a larger radius because it has fewer protons to draw electrons inward. Electronegativity and electron affinity increase to the right across a period and up a group.
The second ionization energy of a neutral atom is always greater than the first.
Recall that ionization energy refers to the energy required to remove an electron from an atom. The first ionization energy will result in a cation. The reduced number of electrons in a cation allows them to be pulled closer to the nucleus by the positively charged protons, effectively decreasing the atomic radius and increasing the attractive force between the electrons and the nucleus. The energy required to remove a second electron must overcome this increased affinity, and will thus be greater than the first ionization energy.
The rest of the answers can be examined with a few simple trends kept in mind. Atomic radius increases to the left of a period and down a group. If two ions have the same electron configuration (like the chlorine anion and potassium cation), then the anion will have a larger radius because it has fewer protons to draw electrons inward. Electronegativity and electron affinity increase to the right across a period and up a group.
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When looking at the periodic table of elements, what is the general trend for increasing electronegativity?
When looking at the periodic table of elements, what is the general trend for increasing electronegativity?
Electronegativity increases to the right and up when looking at the periodic table of elements, such that fourine is the most electronegative element. Following this general trend can help you determine the relative electronegativity between atoms within problems. Remember that periods run horizontally on the table and groups run vertically.
Electronegativity increases to the right and up when looking at the periodic table of elements, such that fourine is the most electronegative element. Following this general trend can help you determine the relative electronegativity between atoms within problems. Remember that periods run horizontally on the table and groups run vertically.
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For any given chemical reaction, one can draw an energy diagram. Energy diagrams depict the energy levels of the different steps in a reaction, while also indicating the net change in energy and giving clues to relative reaction rate.
Below, a reaction diagram is shown for a reaction that a scientist is studying in a lab. A student began the reaction the evening before, but the scientist is unsure as to the type of the reaction. He cannot find the student’s notes, except for the reaction diagram below.
Using NMR, the scientist in the passage determines that there is a negative halide ion present in the products when the reaction reaches step 5. Which of the following factors would tend to raise the energy level at step 5?
For any given chemical reaction, one can draw an energy diagram. Energy diagrams depict the energy levels of the different steps in a reaction, while also indicating the net change in energy and giving clues to relative reaction rate.
Below, a reaction diagram is shown for a reaction that a scientist is studying in a lab. A student began the reaction the evening before, but the scientist is unsure as to the type of the reaction. He cannot find the student’s notes, except for the reaction diagram below.
Using NMR, the scientist in the passage determines that there is a negative halide ion present in the products when the reaction reaches step 5. Which of the following factors would tend to raise the energy level at step 5?
If the halide was flouride, it would have a negative charge dispersed over a smaller cross sectional area than if it was any of the other halides. Flouride is the smallest possible halide, and thus has the highest energy because it concentrates negative charge on the smallest area.
If the halide was flouride, it would have a negative charge dispersed over a smaller cross sectional area than if it was any of the other halides. Flouride is the smallest possible halide, and thus has the highest energy because it concentrates negative charge on the smallest area.
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Place the following atoms in decreasing order of atomic radius.
Cl, Ar, K, Ca
Place the following atoms in decreasing order of atomic radius.
Cl, Ar, K, Ca
Atomic radius has two general trends which you should remember:
1. Atomic radius will decrease when moving left to right along a period.
2. Atomic radius will increase when moving down a group.
Since potassium (K) and calcium (Ca) are farther down the group than argon (Ar) and chlorine (Cl), we conclude that they are the largest atoms in the set. Because Ca is to the right of K, it is slightly smaller than K. As a result, K has the largest atomic radius in the set followed by Ca. Since Ar is to the right of Cl, Cl has a larger atomic radius than Ar.
The decreasing order is K, Ca, Cl, Ar.
Atomic radius has two general trends which you should remember:
1. Atomic radius will decrease when moving left to right along a period.
2. Atomic radius will increase when moving down a group.
Since potassium (K) and calcium (Ca) are farther down the group than argon (Ar) and chlorine (Cl), we conclude that they are the largest atoms in the set. Because Ca is to the right of K, it is slightly smaller than K. As a result, K has the largest atomic radius in the set followed by Ca. Since Ar is to the right of Cl, Cl has a larger atomic radius than Ar.
The decreasing order is K, Ca, Cl, Ar.
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Rank the following ions in order of decreasing atomic radius.
Rank the following ions in order of decreasing atomic radius.
First, you should see that all four of these ions have the same amount of electrons, resulting in a stable \[Kr\] electron shell. These ions differ, however, by the number of protons in their nuclei. Since the heavier ions have more protons pulling on the same amount of electrons, the atomic radius will be smaller, as the negative electrons are drawn inward toward the positive nucleus.
Sr2+ has the most protons in its nucleus out of this set, so it will have the smallest atomic radius. The atomic radius will increase with decreasing atomic number.
The correct order from largest to smallest is Se2-, Br-, Rb+, and Sr2+.
First, you should see that all four of these ions have the same amount of electrons, resulting in a stable \[Kr\] electron shell. These ions differ, however, by the number of protons in their nuclei. Since the heavier ions have more protons pulling on the same amount of electrons, the atomic radius will be smaller, as the negative electrons are drawn inward toward the positive nucleus.
Sr2+ has the most protons in its nucleus out of this set, so it will have the smallest atomic radius. The atomic radius will increase with decreasing atomic number.
The correct order from largest to smallest is Se2-, Br-, Rb+, and Sr2+.
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Which of the following elements has the highest electronegativity?
Which of the following elements has the highest electronegativity?
Remember that electronegativity increases as you approach the top right corner of the periodic table. Since oxygen is the farthest right and the highest up on the perioidic table out of these choices, we conclude that it has the highest electronegativity.
Remember that electronegativity increases as you approach the top right corner of the periodic table. Since oxygen is the farthest right and the highest up on the perioidic table out of these choices, we conclude that it has the highest electronegativity.
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Which of the following elements has the greatest effective nuclear charge?
Which of the following elements has the greatest effective nuclear charge?
The effective nuclear charge is the attractive charge a valence electron feels from the nucleus. As you move from left to right along a period, and more positive charges (protons) fill up the nucleus, the more attraction the valence electron feels. As you move down a group, you jump into the next electron shell, thus shielding the valence electrons from the inner positive charge, and decreasing the effective nuclear charge.
Because chlorine is in the same period as phosphorus and sodium, but has the most protons in its shell (the most right within the same period) it has the greatest effective nuclear charge. Additionally, because chlorine is in the same group as bromine, but is higher up on the periodic table, it has a greater effective nuclear charge, making it the correct answer.
The effective nuclear charge is the attractive charge a valence electron feels from the nucleus. As you move from left to right along a period, and more positive charges (protons) fill up the nucleus, the more attraction the valence electron feels. As you move down a group, you jump into the next electron shell, thus shielding the valence electrons from the inner positive charge, and decreasing the effective nuclear charge.
Because chlorine is in the same period as phosphorus and sodium, but has the most protons in its shell (the most right within the same period) it has the greatest effective nuclear charge. Additionally, because chlorine is in the same group as bromine, but is higher up on the periodic table, it has a greater effective nuclear charge, making it the correct answer.
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Which element would you expect to have the greatest second ionization energy?
Which element would you expect to have the greatest second ionization energy?
Ionization energy is the energy required by an element to remove a valence electron and gain a positive charge. Because alkali metals and alkaline earth metals have one and two valence electrons, respectively, their first ionization energies are both relatively low; shedding a valence electron and absorbing a positive charge actually stabilizes these elements.
Alkaline earth metals can easily release their second valence electron, and have relatively low second ionization energies. Removing this second electron gives these elements a noble gas electronic configuration. This is why the ions of alkaline earth metals are most commonly Mg2+ and Ca2+. In comparison, alkali metals cannot easily shed a second valence electron because it would require removing an electron from an already-filled valence shell. Because of this, alkali metals have extremely HIGH second ionization energies in comparison to alkaline earth metals.
Of the answer choices, potassium would have the highest second ionization energy because it is an alkali metal, rather than an alkaline earth metal.
Ionization energy is the energy required by an element to remove a valence electron and gain a positive charge. Because alkali metals and alkaline earth metals have one and two valence electrons, respectively, their first ionization energies are both relatively low; shedding a valence electron and absorbing a positive charge actually stabilizes these elements.
Alkaline earth metals can easily release their second valence electron, and have relatively low second ionization energies. Removing this second electron gives these elements a noble gas electronic configuration. This is why the ions of alkaline earth metals are most commonly Mg2+ and Ca2+. In comparison, alkali metals cannot easily shed a second valence electron because it would require removing an electron from an already-filled valence shell. Because of this, alkali metals have extremely HIGH second ionization energies in comparison to alkaline earth metals.
Of the answer choices, potassium would have the highest second ionization energy because it is an alkali metal, rather than an alkaline earth metal.
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Which of the given atoms has the largest atomic radius?
Which of the given atoms has the largest atomic radius?
Lithium, boron, oxygen, and neon are all in the same row (period) of the periodic table.
The atomic radius decreases from left to right along a period due to increased effective nuclear force. From left to right the atomic number increases, indicating that more protons are added. The addition of protons increases the positive charge in the nucleus, pulling in the outer electrons by increasing the effective nuclear force, decreasing the radius.
In math terms, we can equate effective nuclear force using the force equation between two charged particles.
We can see that the farther apart the electrons and protons are, the less the force is between them.
Lithium, boron, oxygen, and neon are all in the same row (period) of the periodic table.
The atomic radius decreases from left to right along a period due to increased effective nuclear force. From left to right the atomic number increases, indicating that more protons are added. The addition of protons increases the positive charge in the nucleus, pulling in the outer electrons by increasing the effective nuclear force, decreasing the radius.
In math terms, we can equate effective nuclear force using the force equation between two charged particles.
We can see that the farther apart the electrons and protons are, the less the force is between them.
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Which of the given atoms has the smallest atomic radius?
Which of the given atoms has the smallest atomic radius?
Nitrogen, phosphorous, antimony, and bismuth are all in the same group (column) of the periodic table.
The atomic radius increases from the top of a group to the bottom, due to increased principle shell number (n). As one travels down a group, another s shell is added, meaning that electrons are added in another orbit farther from the nucleus. This serves to increase the atomic radius of the atom.
Nitrogen, phosphorous, antimony, and bismuth are all in the same group (column) of the periodic table.
The atomic radius increases from the top of a group to the bottom, due to increased principle shell number (n). As one travels down a group, another s shell is added, meaning that electrons are added in another orbit farther from the nucleus. This serves to increase the atomic radius of the atom.
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Which of the given atoms has the greatest metallic character?
Which of the given atoms has the greatest metallic character?
Lithium, potassium, rubidium, and francium are all alkali metals in the same group of the periodic table.
Metallic character increases as one moves down each group the periodic table. This is generally due to an increase in the distance between the electrons and the nuclear protons, making it easier for an electron to be lost to a bond or to move freely during electric conduction. Remember that metallic characteristics include malleability, ductility, luster, and shininess.
Lithium, potassium, rubidium, and francium are all alkali metals in the same group of the periodic table.
Metallic character increases as one moves down each group the periodic table. This is generally due to an increase in the distance between the electrons and the nuclear protons, making it easier for an electron to be lost to a bond or to move freely during electric conduction. Remember that metallic characteristics include malleability, ductility, luster, and shininess.
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Which of the given atoms has the least metallic character?
Which of the given atoms has the least metallic character?
Rubidium, ruthenium, tin, and xenon are all in the same row (period) of the periodic table.
Metallic character decreases as one moves from left to right across a period. This is generally due to a reduction in the number of electrons in the outermost shell being from the s orbital. Increased numbers of protons reduce the atomic radius, making it harder for outer shells to have electron motility. Metals will generally have larger radii, allowing for greater electron freedom, ionization, and conductivity. Remember that metallic characteristics include malleability, ductility, luster, and shininess.
Rubidium, ruthenium, tin, and xenon are all in the same row (period) of the periodic table.
Metallic character decreases as one moves from left to right across a period. This is generally due to a reduction in the number of electrons in the outermost shell being from the s orbital. Increased numbers of protons reduce the atomic radius, making it harder for outer shells to have electron motility. Metals will generally have larger radii, allowing for greater electron freedom, ionization, and conductivity. Remember that metallic characteristics include malleability, ductility, luster, and shininess.
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Which of the given atoms has the greatest ionization energy?
Which of the given atoms has the greatest ionization energy?
Fluorine, chlorine, bromine, and iodine are all halogens in the same group of the periodic table.
The ionization energy of an atom, defined as the energy required to remove an electron from an atom’s orbit, decreases as one moves down a group of the periodic table. This is primarily because the atoms increase in size as one moves down the group, thus decreasing the force between the nucleus and an electron in an outermost orbit. It requires less energy to remove an electron that is farther from the protons, than to remove one that is held closer to the nucleus.
Fluorine, chlorine, bromine, and iodine are all halogens in the same group of the periodic table.
The ionization energy of an atom, defined as the energy required to remove an electron from an atom’s orbit, decreases as one moves down a group of the periodic table. This is primarily because the atoms increase in size as one moves down the group, thus decreasing the force between the nucleus and an electron in an outermost orbit. It requires less energy to remove an electron that is farther from the protons, than to remove one that is held closer to the nucleus.
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Which of the given atoms has the lowest ionization energy?
Which of the given atoms has the lowest ionization energy?
Cesium, tungsten, platinum, and radon are all in the same row (period) of the periodic table.
The ionization energy of an atom, defined as the energy required to remove an electron from an atom’s orbit, increases as one move from left to right across the periodic table. This is primarily because the atoms decrease in atomic size, meaning that the outermost electrons that would be removed to create an ion are closer to the nucleus and require more work to remove. Cesium has the largest atomic radius and attains an octet if an electron is removed, making it more stable as an ion, meaning that it will be fairly easy to remove an electron from cesium compared to the other options.
Cesium, tungsten, platinum, and radon are all in the same row (period) of the periodic table.
The ionization energy of an atom, defined as the energy required to remove an electron from an atom’s orbit, increases as one move from left to right across the periodic table. This is primarily because the atoms decrease in atomic size, meaning that the outermost electrons that would be removed to create an ion are closer to the nucleus and require more work to remove. Cesium has the largest atomic radius and attains an octet if an electron is removed, making it more stable as an ion, meaning that it will be fairly easy to remove an electron from cesium compared to the other options.
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