Chemistry Chapter 6 Study Guide

Which particle has the larger radius in each atom/ion pair? a. Na, Na+ b. S, S2− c. I, I− d. Al, Al3+

A. Na B. S^2- C. I- D. Al

Which equation represents the first ionization of an alkali metal atom? a. b. c. d.

C (one with potassium)

What trend is demonstrated by the following series of equations?

Electronegativity

Which noble gas does not have eight electrons in its highest occupied energy level?

Helium, because it is still stable due to the highest energy level being filled.

In general, how are metalloids different from metals and nonmetals?

The only substantial difference is that, under different conditions, metalloids can metals and/or nonmetals.

What pattern is revealed when the elements are arranged in a periodic table in order of increasing atomic number?

The periodic law is revealed, which says that when elements are arranged in order of increasing atomic number, there's a periodic repetition of their physical and chemical properties. Elements of similar properties are grouped together; atoms decrease left to right, (left to right) elements group from metals to metalloids to nonmetals, and similar electron configurations aregrouped together.

Define ion

an atom or group of atoms that has a positive or negative charge

What are the three classes of elements?

metals, nonmetals, and metalloids

What group has multiple charges and are usually cations?

transition metals

Define Inner Transition metal

an element in the lanthenide or actinide series; the highest occupied s sublevel and nearby f sublevel of its atoms generally contain electrons; also called inner transition element

How do positive and negative ions form?

when electrons are transferred between atoms

Define nonmetal

an element that tends to be a poor conductor of hear and electric current; nonmetals generally have properties opposite to those of metals

Define cation

any atom or group of atoms with a positive charge

Define anion

any atoms or groups of atoms with a negative charge

Define Alkali metal

any metal in Group 1A of the periodic table

Define Alkaline Earth metal

any metal in Group 2A of the periodic table

Define Electronegativity

the ability of an atom to attract electrons when the atom is in a compound

What is usually displayed in the periodic table?

the symbols, names of elements, and information about the structure of their atoms

Define Representative element

an element in an "A" group in the periodic table; as a group these elements display a wide range of physical and chemical properties in their atoms, the s and p sublevels in the highest occupied energy level are partially filled

List the symbols for all the elements with electron configurations that end as follows. Note: Each n represents an energy level. a. ns1 b. ns2np4 c. ns2nd10

A. H, Li, Na, K, Rb, Cs, Fr B. O, S, Se, Te, Po C. Zn, Cd, Hg, Uub

Define periodic law

when the elements are arranged in order of increasing atomic number, theer's a periodic repetition of their physical and chemical properties

What are the general trends of ionic sizes?

tends to increase from top to bottom within a group.

How did early chemists sort elements into groups?

by using the elements' properties

What are the 4 physical & chemical properties of Group 8A (noble gases)?

1. All elements are monatomic gases at STP. 2. Noble gases are colorless and odorless. 3. The first compound of a noble gas, XePtF^6, was made in 1962. More than 100 compounds of fluorine and xenon are now known. 4. A compound of argon, HArF, exists only at temp. below -246 C.

What are the 2 atomic properties of Transition Metals?

1. Among them, as atomic number increases, there's an increase in the number of electrons in the second-to-highest occupied energy level. 2. In Periods 5 and 6, transition metals in the same group have identical or almost identical atomic radii. Thus, these pairs of elements have very similar chemical properties. They tend to occur together in nature and are difficult to separate.

Explain Why does it take more energy to remove a 4s electron from zinc than from calcium?

I believe the 4s orbital is the same distance from the nucleus for all atoms. Zinc has a greater positive nuclear charge than calcium because there are more protons. The attraction from the nucleus is greater in the zinc's 4s orbital because of this, causing it to be difficult for the electrons to escape its pull. This is what causes the greater ionization energy.

What electrons are contained in the valence configuration?

all the electrons in a principal energy level

Define noble gas

an element in Group 8A of the periodic table; the s and p sublevels of the highest occupied energy level are filled

Why does taking electrons form noble gases cost energy?

because the energy level is occupied and stable, so it wants to stay stable

How did Mendeleev arrange the elements in his periodic table?

by order of increasing atomic mass

How are the elements arranged in the modern periodic table?

by order of increasing atomic number

What are the general trends of cation and anion (ions) sizes?

decreases from left to right across a period

What do nonmetals do to become anions? (is opposite for metals)

gains electrons (opposite for metals)

What are the two liquid metals?

gallium and mercury

How do you write the Valence Configuration? (tell me if i'm wrong or need more detail)

group label # tells electrons in valence, period # tells the principal energy level, and the block letter tells which orbital it is

In general, what is the trend of atomic sizes?

increases from top to bottom within a group and decreases from left to right across a period

Memorize elements 1-20 on the Periodic Table

know their positions so that you know how they relate to the trends

What is the most electronegative atom? The least electronegative? What Value?

most=fluorine=4.0 lowest=cesium

Why did Mendeleev leave spaces in his periodic table?

Mendeleev left these spaces in his table because he knew that bromine belonged with chlorine and iodine. He predicted that elements would be discovered to fill those spaces, and he predicted what their properties would be based on their locations in the table.

Make Generalizations Why is the first ionization energy of a nonmetal much higher than that of an alkali metal?

Nonmetals generally tend to be smaller than metals (such as alkali metals). This indicates that there are less occupied orbitals meaning the ones that are occupied are closer to the nucleus. The electrons therefore receive a greater attraction from the nucleus because of the closer nucleus. This stronger attraction causes the ionization energies for nonmetals to be greater than that of the alkali metals. In other words, nonmetals are so close to achieving the same number of electrons as a noble gas, do it really doesn't want to give up any more electrons.

Which of the following are symbols for representative elements: Na, Mg, Fe, Ni, Cl?

Sodium, Magnesium, and Chlorine

What effect did the discovery of gallium have on the acceptance of Mendeleev's table?

The discovery of gallium and germanium helped convince scientists that Mendeleev's periodic table was a powerful tool because Mendeleev's predicted properties of these missing elements was such a close march to their actual properties.

How does the ionic radius of a typical metal compare with its atomic radius?

The ionic radius of a metal cation is smaller than the atomic radius of the metal atom. Metals tend to form cations. Cations usually are smaller than their original atom because of the overpowering positive charge of the nucleus attracting the negative electrons, making the size/radius smaller. Thus, the ionic radii of a typical metal are small when compared to its atomic radius.

Why are noble gases not included in the Electronegativity Values for Selected Elements figure?

The noble gases aren't included because they don't usually form compounds. Thus, there's not really any good reason for the electronegativity values to be listed since the values don't occur often.

Define Halogen

a nonmetal in Group 7A of the periodic table

Define metalloid

an element that tends to have properties that are similar to those of metals and nometals

What are the general trends of electronegativity values?

decreases from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period.

What are the names of the elemental rows in the f block?

lyntenides (top, not sure if this is the right name) actinide (bottom)

What are the elements that are out of the transition metal blocks?

main or representative elements

What can elements be sorted into based on their electron configurations?

noble gases, representative elements, transition metals, or inner transition metals

Define metal

one of a class of elements that are good conductors of heat and electric current; metals tend to be ductile, malleable, and shiny

Define Transition metal

one of the Group B elements in which the highest occupied s sublevel and a nearby d sublevel generally contain electrons

Define atomic radius

one-half the distance between the model of two atoms of the same element when the atoms are joined

What are the general trends of first ionization energy?

tends to decrease from top to bottom within a group and increases from left to right across a period

Define ionization energy

the energy required to remove an electron from an atom in its gaseous state

Which element in each pair has a higher electronegativity value? a. Cl, F b. C, N c. Mg, Ne d. As, Ca

(electronegativity increases across period, decreases down a group) A. Fluorine (the most electronegative) B. Nitrogen C. Magnesium (neon doesn't form compounds) D. Calcium

Arrange the following groups of elements in order of increasing ionization energy: a. Be, Mg, Sr b. Bi, Cs, Ba c. Na, Al, S

(increases across period, decreases down a group) A. Sr, Mg, Be B. Cs, Ba, Bi C. Na, Al, S

Write the electron configurations of these elements. a. the noble gas in Period 3 b. the metalloid in Period 3 c. the alkali earth metal in Period 3

(notice that all have their highest occupied in principal energy level 3) A. Argon, 1s^2 2s^2 2p^6 3s^2 3p^6 B. Silicon, 1s^2 2s^2 2p^6 3s^2 3p^2 C. Magnesium, 1s^2 2s^2 2p^6 3s^2

*Using a diagram and words fully explain how different metallic ions produce their unique color "barcode" (atomic emission spectrum)*(part 2)

(use hydrogen atomic emission spectrum) This represents an element/atom's atomic emission spectrum. The lines labeled "ground state" represent the state in which electrons aren't excited. The "n-2" lines represent the light emitted from excited electrons that drop down to the 2nd energy level. The last section represents the light emitted from excited electrons falling back to energy levels higher than n=2.

*Using a diagram and words fully explain how different metallic ions produce their unique color "barcode" (atomic emission spectrum)*(part 1)

(use the diagram from our flame test lab sheet) When metallic ions absorb energy usually in the form of heat, electrons are raised from lower energy levels to higher energy levels. In this state, electrons are excited from the energy absorption. This is unstable, however, so the electrons alternatively drop down to lower energy levels, emitting light with a wavelength equal to the energy absorbed.

What are the 3 physical properties of Group 3A? (group is immediately after the B/d block)

1. B is a metalloid. The rest of Group 3A are metals. 2. Al is a valuable structural material because of its strength, especially in alloys with silicon or iron. These alloys have a low density and resist corrosion. 3. Gallium has an extremely wide liquid temp. range (30 C to 2204 C). Solid gallium floats in liquid gallium, which is unusual in a metal.

For which of these properties does lithium have a larger value than potassium? a. first ionization energy b. atomic radius c. electronegativity d. ionic radius

A and C

Which of these metals isn't a transition metal? a. aluminum b. silver c. iron d. zirconium

A, Aluminum

Which element in each pair has a greater first ionization energy? a. lithium, boron b. magnesium, strontium c. cesium, aluminum

A. Boron B. Magnesium C. Aluminum

Identify each property below as more characteristic of a metal or a nonmetal. a. a gas at room temperature b. brittle c. malleable d. poor conductor of electric current e. shiny

A. Nonmetal B. Nonmetal C. Metal D. Nonmetal E. metal

When the elements in each pair are chemically combined, which element in each pair has a greater attraction for electrons? a. Ca or O b. O or F c. H or O d. K or S

A. Oxygen B. Fluorine (the most electronegative) C. Oxygen D. Sulfur

Which element in each pair has atoms with a larger atomic radius? a. sodium, lithium b. strontium, magnesium c. carbon, germanium d. selenium, oxygen

A. Sodium B. Strontium C. Germanium D. Selenium

Explain why there should be a connection between an element's electron configuration and its location on the periodic table.

An electron's electron configuration determines its location (group) in the periodic table. Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals based on electron configurations. Electrons play a key role in determining the properties of elements so there should be a connection between an element's electron configuration and its location in the periodic table. The period number determines the principal energy level, the letter of the blocks tell the orbital, and the group number gives the number of electrons occupying that orbital group.

In which pair of elements are the chemical properties of the elements most similar? Explain your reasoning. a. sodium and chlorine b. nitrogen and phosphorus c. boron and oxygen

B., Nitrogen and Phosphorus are within the same group (Group 5A), which indicates that they share similar chemical properties.

*Using a diagram and words fully explain how different metallic ions produce their unique color "barcode" (atomic emission spectrum)*(part 3)

Electrons produce unique barcodes of light. Depending on their number of electrons and their electron configurations, the "barcode lines" are different because of the varying wavelengths of light an atom emits when absorbing different amounts of energy. (Light isn't produced from a cooling atom, and the excited electrons alternate between lower levels which don't always have to be ground state)

Based on their locations in the periodic table, would you expect carbon and silicon to have similar properties? Explain your answer.

I would expect these two elements to have similar chemical and physical properties because they're located within the same group.

Describe shielding

Each occupied energy level forms a force field. Innermost electrons feel the nucleus' effects the most. More shields means more protection from nucleus. The further away from the nucleus, the more loosely they're held

What does the Octet rule say?

noble gases have 8 electrons in their valence shell, which is what all ions try to achieve because it represents stability

Explain the atomic size trend What about ionic size?

smaller atoms have a stronger attraction of electrons, while the size increases down the group because of the principal energy level. Same reasoning/trend for ionic size

Use the graph to estimate the atomic radius of an indium atom.

About 166 pm

What are the 3 atomic properties of Group 8A (noble gases)?

1. Noble gases have an electron configuration that ends in ns^2 np^6, except for helium (1s^2). 2. In noble gas compounds, the most common oxidation # for the gas is +2. 3. Noble gases have the highest ionization energies because their energy levels are filled.

What are the 4 physical properties of Group 5A?

1. Except for nitrogen gas, they're solid at room temp.. 2. The metallic properties of them increase from top to bottom within the group. N and P are nonmetals. As and Sb are metalloids. Bi is a metal. 3. Liquid N is a cryogen, a liquid refrigerant that boils below -190 C. 4. P has 10 allotropes including white and red.

What are the 3 physical properties of Group 6A?

1. Except for oxygen gas, O^2, they're solid at room temp. 2. The metallic properties of them increase from top to bottom within the group. 3. Polonium is a radioactive metal.

What are the three physical properties of the Transition Metals? (the B/d block)

1. Most are ductile, malleable, and good conductors of heat and electric current. 2. For them, density tends to increase across a period, while melting point increases to a peak in Group 6B and then decreases. 3. Compounds tend to have color.

What are the 4 atomic properties of Group 3A?

1. The elements have an electron configuration that end in ns^2 np^1 2. The most common oxidation # for B, Al, gallium, and Indium is +3. For thallium, it's +1 3. They become more metallic from top to bottom within the group. 4. Radioactive thallium-201 is injected into patients taking a stress test used in the diagnosis of heart disease.

What are the 3 chemical properties of Transition Metals.?

1. There's great variation in reactivity among transition metals. Scandium and yttrium are similar to Groups 1A and 2A metals. They're easily oxidized on exposure to air and react with water to release hydrogen. Platinum and gold are extremely unreactive and resist oxidation. 2. In general, transition metals have multiple oxidation states. Compounds in which these elements are in their highest oxidation stares are powerful oxidizing agents. 3. Most form compounds with distinctive colors. The color of a transition metal compound or solution can indicate the oxidation state of the metal.

What are the 4 atomic properties of Group 7A?

1. They have an electron configuration ending in ns^2 np^5. 2. They exist as diatomic molecules. 3. Each has the highest electronegativity in its period. 4. The most common ionic charge for them is 1-. Except for fluorine, halogens have positive oxidation #s of +1, +3, +5, and +7.

What are the 3 atomic properties of Group 4A?

1. They have an electron configuration that ends in ns^2 np^2. 2. For these, the most common oxidation #s are +4 and +2. For carbon, -4 is also common. 3. Silicon and germanium are semiconductors.

What are the 4 atomic properties of Group 5A?

1. They have an electron configuration that ends in ns^2 np^3. 2. The most common oxidation #s for them are +3, +5, and -3 3. N has oxidation numbers from -3 to +5 (and all #s in between) in a variety of stable compounds. 4. Elemental nitrogen, N^2, is a highly unreactive due to its strong N-to-N triple bond.

What are the 3 atomic properties of Group 6A?

1. They have electron configurations that end in ns^2 np^4. 2. For them, most common oxidation #s are +4, +6, and -2. 3. Oxygen is paramagnetic because there are unpaired electrons in O^2 molecules.

What are the 3 physical properties of 7A (Halogens)?

1. They're nonmetals. At room temp., fluorine and chlorine are gases and bromine is a liquid. Iodine and astatine are solids. 2. They're very reactive. The reactivity decreases from fluorine to astatine. They don't exist in the elemental form in nature. 3. Astatine isotopes are radioactive with short half-lives.

What are the 3 physical properties for Group 4A?

1. all solids at room temp. 2. Their metallic properties increase from carbon to lead. 3. Diamond, graphite, and buckminsterfullerene (who came up with this name?!) are 3 allotropes of carbon

What are the 5 atomic properties of Group 1A (Alkali Metals)?

1. an electron configuration that ends in ns^1 2. the most reactive metals (=low ionization energies) 3. form ions with a 1+ charge 4. atoms are the largest of their periods 5. Cesium is a good reducing agent because its first ionization energy is very low

What are the 5 atomic and physical properties of Hydrogen?

1. has an electron configuration of 1s^1. 2. Most common oxidation #s for it are +1 and -1. 3. Most hydrogen (99.985%) is protium, or hydrogen-1. 4. The other stable isotope is deuterium (hydrogen-2), which has the symbol D. Harold Urey discovered heavy hydrogen, D^2, in 1931. 5. Tritium (hydrogen-3) was discovered in 1934. Its half-life is 12.3 years.

What are the 4 atomic properties of Group 2A (Alkaline Earth Metals)?

1. have an electron configuration that ends in ns^2. 2. strong reducing agents, losing 2 electrons and forming ions with a 2+ charge. 3. Because radium is luminous, it was once used to make the hands and numbers on watches glow in the dark. 3. The ratio of ^87 Sr to ^86 Sr varies with location. This data is used to solve puzzles, such as the source of timber used in prehistoric buildings.

What are the 4 physical properties of Group 2A (Alkaline Earth Metals)?

1. relatively soft, but harder than alkali metals. 2. gray-white luster when freshly cut. When exposed to air, they quickly form a tough, thin oxide coating. 3. Densities, melting pints, and boiling points tend to be higher than for the alkali metal in the same period. 4. Magnesium alloys are strong and lightweight. They're used in cameras, lawnmowers, aircraft, and automobiles.

What are the two main physical properties of Group !A (Alkali Metals)?

1. silver-gray solids that are soft enough to cut with a knife. They're soft because they have only one valence electron. 2. Presence of a single valence electron explains the low melting and boiling points of alkali metals.

The bar graph shows how many elements were discovered before 1750 and in each 50-year period between 1750 and 2000. a. In which 50-year period were the most elements discovered? b. How did Mendeleev's work contribute to the discovery of elements? c. What percent of these elements were discovered by 1900?

A. 1801-1850 B. His table arranged elements in a way more accepted than other tables. Scientists were able to better see the trends of the elements. This probably helped scientists guess where some elements' locations were on this table and what properties they would have. His table also led to the development of the modern periodic table, which probably helped a great deal in discovering elements. C. 75.2%

Apply Concepts Write the electron configurations of the following ions: a. the liquid in Group 7A with a 1− charge b. the metalloid in Period 3 with a 4+ charge c. the gas in Group 6A with a 2− charge d. the alkali earth metal in Period 3 with a 2+ charge

A. Bromine, 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 B. Silicon, 1s^2 2s^2 2p^6 C. Oxygen, 1s^2 2s^2 2p^2 D. Magnesium, 1s^2 2s^2 2p^6

Write the symbol of the element or elements that fit each description. a. a nonmetal in Group 4A b. the inner transition metal with the lowest atomic number c. all of the nonmetals for which the atomic number is a multiple of five d. a metal in Group 5A

A. C (within group, trend going down a group is that it becomes more metallic) B. La C. P and Br D. Bi

In each pair, which ion is larger? a. Ca2+, Mg2+ b. Cl−, P3− c. Cu+, Cu2+

A. Calcium B. P3− C. Cu+

Locate each of the following elements in the periodic table and decide whether its atoms are likely to form anions or cations. a. sodium b. fluorine c. calcium d. potassium e. iodine f. beryllium g. oxygen h. lithium

A. Cations B. Anions C. Cations D. Cations E. Anions F. Cations G. Anions H. Cation

Interpret Diagrams Use the periodic table and the Electron Configurations figure to identify the following elements: a. has its outermost electron in 7s1 b. contains only one electron in a d orbital

A. Francium b. Scandium

The bar graph shows the relationship between atomic and ionic radii for Group 1A elements. a. Describe and explain the trend in atomic radius within the group. b. Explain the difference between the size of the atoms and the size of the ions.

A. The atomic radius increases (in a group) with the increasing atomic number. This is because atomic size generally increases from top to bottom within a group because the greater the number of protons means the greater amount of electrons. The amount of electrons in orbitals determine an atom's size/radius. B. Since Group 1A (Alkali Metals) are metals, they tend to form cations. Cations are always smaller than their original atom because the greater positive charge from the nucleus closes in the space between it and the electrons, thus "shrinking" its size. This is why the ionic radii are smaller than the atomic radii of the same element.

Explain Give a reason for each of the following comparisons: a. Calcium has a smaller second ionization energy than does potassium. b. Lithium has a larger first ionization energy than does cesium. c. Magnesium has a larger third ionization energy than does aluminum.

A. The electrons in calcium are removed from the same energy level. In potassium, the second electron is removed from a lower energy level. Removing two electrons from Potassium would require more energy because that would mean taking electrons from a different energy level, while removing two from calcium doesn't cost as much energy because it wants its energy levels to be completely occupied. Also, calcium has a higher atomic number, and therefore size, than potassium. This means that the electrons are further from the nucleus' pull. The weaker pull allows electrons to be removed easier. B. Because cesium has a larger atomic radius than lithium, the nuclear charge in a cesium atom has a smaller effect on the electrons in the highest occupied energy level. C. The third electron removed from a magnesium atom is in a different,lower energy level that therefore requires more energy. The third electron removed from aluminum, however, empties out its principal energy level so that now all its energy levels are fully occupied, which is what it wants.

Interpret Graphs The graphs show the relationship between the electronegativities and first ionization energies for Period 2 and Period 3 elements. a. Based on data for these two periods, what is the general trend between these two values? b. Use nuclear charge and shielding effect to explain this trend.

A. These two values have a positive relationship, with electronegativity increasing with the first ionization energies of these elements. B. The smaller an atom is, the greater the attraction of the positive nuclear charge will be on the electrons surrounding it. The ionization energies are greater the smaller the atom is because it is more difficult for electrons to be removed. The bigger an atom is the more orbitals there are to shield the electrons from this nuclear charge, overcoming the attraction and allowing the atom to increase in size. The electronegativity corresponds to the ionization energies because atoms with a greater nuclear charge not shielded by more orbitals attracts more electrons in a compound.

Explain why fluorine has a smaller atomic radius than both oxygen and chlorine.

Fluorine has a smaller atomic radius than both oxygen and chlorine because chlorine has electrons in a higher principal energy level, so the radius is smaller in comparison. Oxygen, however, has less protons to pull electrons in the same principal energy level closer, allowing the radius to not be as small as fluorine (whose nucleus is more positive, and thus shrinks the distance between it and its electrons).

Would you expect metals or nonmetals in the same period to have higher ionization energies? Give a reason for your answer

I would expect the nonmetals sharing a period with metals to have higher ionization energies because the farther right along a period you go the higher the ionization energy. This happens because at the beginning of a period is a new energy level with minimal electrons occupying it. The farther right you go, the more protons are within the nucleus due to the increasing electrons that occupy the energy level of that period. The attraction from the protons to the electrons increases left to right in a period, corresponding with its ionization energy.

There is a large jump between the second and third ionization energies of magnesium. There is a large jump between the third and fourth ionization energies of aluminum. Explain these observations. 61. The bar graph shows the relationship between atomic and ionic radii for Group 1A elements. a. Describe and explain the trend in atomic radius within the group. b. Explain the difference between the size of the atoms and the size of the ions

It's relatively easy to remove two electrons from Magnesium; it's much more difficult to remove a third electron. It's relatively easy to remove three electrons from aluminum; it's much more difficult to remove a fourth electron. There's a large jump between the 2nd and 3rd ionization energies because Magnesium would have to take electrons from a lower principal energy level, going from a stable atom with fully occupied energy levels to an unstable atom, which atoms don't like to do. The same reasoning applies to Aluminum; there's a large jump between the 3rd and 4th ionization energies because it would cause one of its energy levels/orbitals to be only partially filled.

Where are the alkali metals, the alkaline earth metals, the halogens, and the noble gases located in the periodic table?

The alkali metals are located in Group 1A in the farthest left column. The alkaline earth metals are located to the right of the alkali metals, in Group 2A. The halogens are located in Group 7A, left of the rightmost column. The noble gases are located in Group 8A, in the rightmost column.

Why is there a large increase between the first and second ionization energies of the alkali metals?

The electrons are easier to remove the first time around rather than the second. When you remove the first electron, the atom is now stable, with orbitals that are fully occupied; there are no unpaired electrons to unbalance things. The second electron is harder to remove because then the atom's orbitals won't be so perfectly filled; the atom will become unstable once more, which atoms don't like to do.

Explain the difference between the first and second ionization energy of an element.

The first ionization energy represents the amount of energy required for an electron to be first removed from its atom, while the second ionization energy is merely the energy required to remove a second electron from an atom. The second ionization energy is greater because the nucleus' attraction is greater. The difference in energy between these ionization energies indicates what charge this element tends to form ions. So, if there was a large increase from the first to second ionization energy, then the element will tend to form in a 1+ charge because of the difficulty it takes to remove a second electron.

Sequence The following spheres represent Ca, Ca2+, and Mg2+. Which one is which? Explain your reasoning.

The largest sphere (a) is Ca because Ca has the largest atomic size due to its higher atomic number. The next sphere (b) is Ca^2+ because it is a cation, and cations are always smaller than their original atom because of the "shrinking" quality of the positive ion. The smallest sphere is Mg^2+ because its atomic number/size is smaller than Calcium and is made smaller even more so due to it being a cation, which naturally shrinks the size.

Predict Do you think there are more elements left to discover? If so, what is the lowest atomic number a new element could have? Explain.

There may be more elements left to discover; I'm not sure. I guess it just depends on how many electrons and protons nature will allow to come together in one atom without it being unstable. The lowest atomic number a new element could have is119 because there are already elements for atomic numbers 1-118.