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Name: _________________________________________________ Room 262 The Periodic Table is a truly fantastic table of information about all of the elements known to scientists. The table is laid out in 18 vertical groups, that we usually “read” top to bottom. It has 7 horizontal periods that we read “left to right”. The atoms are all metals, or non-metals, while just 7 are also called the metalloids, which mean that they also have some of the “other team’s” properties. Silicon is non-metal, but it can conduct electricity. Antimony is a metal, but when you bang it with a hammer, rather than acting malleable (flattening), it cracks because it’s brittle. Atoms in GROUPS are very similar chemically and physically. In fact, when it comes to bonding, atoms in a group are generally interchangeable. That is shown when looking at group one metals (that all make +1 cations) bonding to group 17 nonmetals (that all make –1 anions). Most common to you is NaCl, sodium chlo-

• ride. All group one metals bond in a 1:1 ratio with group 17 nonmetals: LiCl, KCl,

RbCl, and CsCl. Also, LiBr, NaBr, KBr, RbBr, etc. Atoms in the same period are not very alike at all. On the right side of the table are the most reactive metals. Midtable there are the less reactive metals, followed by metalloids, nonmetal, and even a noble gas. The ONLY commonality that atoms in a period have is that they contain the same number of electron orbitals, which matches the period that they are in (period 2 atoms have 2 orbitals). Within the table there are many “trends” or patterns. Some of the data we will compare is on the Periodic Table itself, some is on Table S, or even must be in- ferred by you thinking hard. Trends run down the tables (group trends) or across the table (period trends). We will examine seven trends in both directions, and im- portantly, why these trends exist. Finally, there are some “exceptions to these trends” or glitches. We need to see them, and grasp the why they happen as well,

Trends of the Periodic Table Basics Trends are patterns that atoms on the periodic table of elements follow. Trends hold true “most” of the time, but there are exceptions, or “blips”, where the trend seems to do the wrong thing. It is important when investigating a particular trend that you examine at least four atoms in a group or period and see what the trend numbers are doing. Choosing just 2 atoms might show you the exception, rather than the trend itself. The seven trends we study in class are these:

• 1. atomic radius (relative size, measured in picometers, pm) - on Table S.
• 2. average weighted atomic mass (measured in amu) - on the Periodic Table.
• 3. net nuclear charge (how positive is the nucleus, related to # of protons) - on the Periodic Table.
• 4. ion size (cations or anions) - to be inferred from data on the Periodic Table.
• 5. electronegativity (relates to bonding) - on Table S.
• 6. 1st Ionization Energy (energy required to change a mole atoms → a mole +1 cations) - on Table S.
• 7. metal property or non-metal property - on the Periodic Table.

Group Trends: the trend that the atoms follow going down any particular group Period Trends: the trend that the atoms follow going across any particular period Atomic Size—Atomic Radius

Reference table S shows us atomic radius, which is the measure of distance from nucleus to outer most electron orbital. The measurement is in picometers (1 x 10-12 meters). Look up the Atomic Radius for each atom in group 1, and in period 4, then state the trends simply. Copy these two statements just below, and make sure that they make sense to you. Really.

• 1. The period trend for atomic radius is increasing.
• 2. The group trend for atomic radius is increasing.

Atom Radius in pm Li Na K Rb K Ca Sc Ti Atoms → Radius in pm

1 2

The group trend for atomic size is INCREASING. That is because each atom that follows going down a group has one more orbital than the atom above it. Three orbitals are larger than two orbitals, four orbitals are larger than three. Fill in these two charts showing electron orbitals for two more groups.

• 3. State the group trend for atomic radius for group 2 (complete sentence)
• 4. State the group trend for atomic radius for group 17 (complete sentence)

The period trend for atomic size is DECREASING. As you go across a period the atoms have the same number of electron orbitals but each adds an extra proton. With more protons “pulling” inward on the same number of electron orbitals— the atoms get smaller and smaller as you go across the table. The smallest atom in any period is the noble gas, because that’s the atom with the MOST number of protons with a given number of electron orbitals. Fill in this chart for two different periods of atoms and their atomic radius measures.

• 5. State the PERIOD TREND for ATOMIC RADIUS (for any period)

Group 2 atoms Electron orbitals Group 17 atoms Electron orbitals Be F Mg Cl Ca Br Sr I Period 2 Li

• 130. pm

Be B C Period 5 Rb Sr Y Zr

Atomic Mass Atomic mass is measured in amu, atomic mass units. The average weighted atomic mass for each atom is listed on the Periodic Table of Elements. Generally speaking the smallest atoms are those with the lowest atomic numbers, and they get heavier as this number increases. Atomic mass is a measure of the number of protons and neutrons in a nucleus, as we accept that the mass of electrons is so small that we disregard it. One neutron = 1 amu. One proton = 1 amu. One electron = zero in high school. The atomic mass is how many protons and neutrons in total that are in the nucleus. All atoms have isotopes, chemically identical atoms with different masses because they have different numbers

• f neutrons. Neutrons are neutral, they don’t really affect the chemistry or properties (other than mass), so all

isotopes for a given atom react the same way. Fill in these tables for ATOMIC MASS (you can round to the nearest whole number of AMU)

• 6. State the PERIOD TREND for atomic mass.
• 7. State the GROUP TREND for atomic mass.

The period trend for atomic mass is also INCREASING, but there are some exceptions (see cobalt - nickel, and then check Argon-Potassium). Exceptions like this are due to the relative numbers of neutrons in isotopes

• f certain atoms. They are exceptions to the trend, they DO NOT BREAK the pattern.

Na Mg Al Si Atoms → Atomic mass in AMU (u) Atom Atomic mass in AMU (u) O S Se Te

Net Nuclear Charge The subatomic particles: electrons, protons, & neutrons all have particular charges. Electrons are negative (-1) and are all located outside the nucleus. Neutrons are neutral (Ø) and even though they are in the nucleus, add NO CHARGE to the nucleus. The protons of the nucleus are positively charged (+1) and are the measure of net nuclear charge. This trend is a measure of how much positive charge is in the nucleus of the atom, which is measured by how many protons, each with a +1 charge, are in a nucleus of an atom. Since each atom has a certain number of protons (the ATOMIC NUMBER), it’s easy enough to count the net nuclear charges. Examples: He has 2 protons and 2 neutrons in the nucleus, this adds to a +2 net nuclear charge Ar has 18 protons and 22 neutrons in the nucleus, this adds to a +18 net nuclear charge. Fill in these tables for NET NUCLEAR CHARGE. State the GROUP TREND FOR NET NUCLEAR CHARGE. State the PERIOD TREND FOR NET NUCLEAR CHARGE. The group trend for net nuclear charge is INCREASING. The period trend for net nuclear charge is INCREASING. There are NO exceptions to this trend. This would require a + sign to be true. Helium has 2 protons, but the net nuclear charge for helium is +2. 2 is not the same as +2. Na Mg Al Si +11 Atoms → Net Nuclear Charge Atom Net Nuclear Charge O +8 S Se Te

Ion Size Ions come in two varieties, cations are atoms that have lost electrons and become positively charged, and are always metals. Anions are atoms that have gained electrons and become net negatively charged, and are always non-metals. Ions form by gaining or losing enough electrons to get that “special” stable, noble gas electron configuration. When an ion forms, it obtains a noble gas electron configuration, which is called being ISOELECTRIC to a noble gas. These ions are not noble gases, they obtain the same electron configuration as a noble gas. When an atom becomes a cation it loses ALL of its valence, or outermost electrons. Group one atoms all lose

• nly one electron. Group 2 atoms all lose 2 electrons as they become +2 cations. Metals always lose all of

their valence electrons, to become isoelectric to a noble gas. State the GROUP TREND for CATION SIZE. Because cations lose a whole outer (valence) orbital, they are always smaller than the atoms they started out

• as. The sodium cation is smaller than the sodium atom. The calcium cation is smaller than the calcium atom.

The aluminum cation is smaller than the aluminum atom. Cations are always quite a bit smaller than the atoms. Fill in this chart of electron configurations for these nonmetals. State the PERIOD TREND FOR ANION SIZE. Atom Electron configuration Cation Electron configuration Li Li+1 Na Na+1 K K+1 Rb Rb+1 Atom N O F Atom electron configuration ANION electron configuration

When a non-metal atom becomes an anion it gains enough electrons to obtain a full outer orbital, so it can be ISOELECTRIC to a noble gas. Non-metals can gain one, two, or even three electrons to fill the outer valence

• rbital. When the atoms gain electrons in the valence orbitals this orbital must stretch a bit to accommodate

this influx of negative charge. The electrons all repel each other, and the extra electron will force all the elec- trons in that orbital a bit further away from each other. Anions are always slightly larger than the atoms that they formed from. Fill in this chart for ANION SIZES in groups. State the GROUP TREND for anion size. Atomic size increases going down a group because each atom lower on the table has more orbitals. Same for cations, although a cation is smaller than its atom, each successive cation has more orbitals than the previous. Cations get smaller going across a period as each cation has the same number of electrons as all the other cations in the period, but they gain electrons across the period, so more protons are pulling on the same number of electrons. Cations Na+1 , Mg+2, and Al+3 all have 10 e- in a 2-8 configuration but have 11, 12, and 13 protons respectively. Sodium is larger than magnesium because Mg has that extra proton pulling the outer

• rbital in. The aluminum +3 ion is smallest because it has yet another proton pulling on the same number of

electrons in the same number of orbitals. The Group Trend for Cations is INCREASING. The Period Trend for Cations is DECREASING. Anions are all larger than their atoms because they squeeze an extra electron into the outer orbital, where it can fit, but forces all the negatively charged electrons a bit further away from themselves. Looking at group 16,

• xygen sulfur and selenium, all get larger as atoms moving down the table. Each anion is larger than it’s atom,

and the anions get progressively larger as well. For a period trend, look at period 3, phosphorous, sulfur and chlorine. The atoms get progressively smaller going across the period. Anions for these three form as P-3, S-2, and Cl-1 respectively. Since there are more and more protons in these ions moving across the period, the anions get smaller moving across the period. The Group Trend for Anions is INCREASING. The Period Trend for Anions is DECREASING. Atom Electron configuration ANION Electron configuration F F-1 Cl Cl-1 Br Br-1 I I-1

Electro negativity

Electronegativity is the tendency of an atom to attract electrons to itself in a bond When atoms make covalent bonds, they “share” electrons. Each atom puts up one electron, and when a pair

• f electrons is shared, a single covalent bond forms. Examples include: H2, Cl2, and Br2. Each shares a single

pair of electrons, each makes a single covalent bond. Since these atoms are identical to each other, and they have the same tendency to attract an electron to itself (they have the same electronegativity values), the bond is a single NON-POLAR covalent bond, neither atom “gets” those electrons more than the other atom does. The H2 “shares” the electrons they bond with evenly, each atom has a 2.2 electronegativity value (from Ta- ble S). The difference between them is zero 2.2—2.2 = 0 The smaller the difference, the more evenly shared the electrons are. This bond is NONPOLAR. With HF, the atom’s electronegativity values are 2.2 and 4.0, a difference of 1.8! That is very polar. Here, although the electrons are “shared”, fluorine pulls MUCH, MUCH harder and “gets” the electron from hydrogen most of the time, filling it’s outer orbital MOST

• f the time. This makes the F atom more negative most of the time, and the H atom more positive MOST of

the time. This is a very POLAR BOND. Electronegativity allows a student to compare bonds, are they nonpolar, or if they are polar, which are the most polar (biggest difference in electronegativity values), or less polar (smaller EN difference). Which of these bonds is MOST POLAR, or LEAST POLAR? Compound HCl HBr H2O HI Electronegativity values 2.2 and 3.4 Difference in EN VALUES 1.2 Which is MOST POLAR, LEAST POLAR?

Fill in these tables with Electronegativity values, and

• 8. State the group trend for Electronegativity.
• 9. State the group trend for Electronegativity.

When H2O, or CO2 forms, The atoms make more than one bond. In water, one oxygen atom bonds to two separate hydrogen atoms. Each bond is independent when it concerns electronegativity. With CO2, a carbon atom makes 2 separate bonds to the oxygen atoms. Here though, the oxygen atom will form a double bond to

• carbon. In CO2, there are two double polar covalent bonds. You know this because the electronegativity

values are 2.6 for carbon and 3.2 for the oxygen atoms. The oxygen atoms pull the electrons in the bond more so to the oxygen side of the bond. This makes the oxygen atoms “more negative”, the carbon atom “more positive”, hence, double POLAR bonds form. ANY difference in electronegativity means the bond is polar. If the difference is ZERO, the bond is non-polar. Na Mg Al Si Atoms → Electronegativity Values Atom Electronegativity Values O S Se Te

Linus Pauling described this concept of electronegativity. He measured that fluorine has the greatest tendency to gain electrons in a bond. Since he created this scale, he could do what he wanted, and he did.

• Dr. Pauling decided that F would have an EN value of 4.0, the highest value on the table. All other atoms

would be compared to Fluorine. Since all atoms are compared to a single “standard” atom, this is a relative

• scale. All atoms are measured, relative or compared to a standard that HE decided upon. Electronegativity

is an example of a relative scale. Electronegativity does not have units.

He choose the numbers 0 to 4.0 for no particular reason. The numbers don’t “mean”

anything, they are just numbers to “rank” the atoms.

Electronegativity is a relative scale, and it is ARBITRARY as well.

Table S shows all the EN values for the elements. Some atoms, the smaller noble gases have no EN values. These atoms do not make bonds ever, they have NO TENDENCY to gain (or lose) electrons. The Group Trend for electro negativity is decreasing. The Period Trend for electro negativity is increasing. The noble gas XENON does have an EN value, and under some unusual conditions it can be forced into a

• bond. This is an exception, noble gases tend to be relatively INERT, or non-reactive.

The period trend for electronegativity is increasing, because in any period you have one orbital, and more and more protons pulling tighter and tighter. There is a greater inward attraction to gain electrons as you move across the table. Until you get to group 18, which already has full orbitals. The group trend for electronegativity is decreasing. This seems odd at first, more and more protons are in the atoms going down any group, but this increased positive charge in the nucleus pulling inwards is offset by the distance the outermost valence orbital is to the nucleus. More protons helps pull inward, distance hurts more than the increased number of protons helps.

The whole table trend is the closer you are to F, the higher the EN value except for Group 18.

1st Ionization Energy

When atoms of group 1 become cations and “lose” an electron, even though they “want” to do this to gain the noble gas electron configuration, it requires some energy. The electrons do not FALL OFF of the atoms—they must be pulled off, and that takes energy. The 1st Ionization energy is the energy required to pull a mole of electrons off of a mole of atoms. The amount of energy required to take a mole of atoms and make them a mole of +1 cations is called the 1st IONIZATION ENERGY. The unit is kJ/mole or kilo-joules per mole.

• 10. State the GROUP TREND FOR 1st Ionization Energy.
• 11. State the PERIOD TREND for 1st Ionization Energy.

The metals have lower first ionization energy requirements for several reasons. Metals will tend to lose electrons easier than nonmetals. Metals form cations. Nonmetals gain electrons to form anions. The largest atoms in any period are in group 1, and these atoms have the lowest 1st ionization energy. Moving across any period, you have the same number of orbitals, but more and more protons, pulling the atoms smaller. For the same reason that the atoms get smaller (greater inward attraction) the more difficult it becomes to pull these electrons off. The group trend for first ionization energy is increasing. Going down any group, since the atoms are getting larger and larger, even though there are more protons, the inward attraction the nucleus on these electrons weakens over this distance. The group trend for 1st ionization energy is decreasing. Na Mg Al Si Atoms → 1st Ionization Energy kJ/mole Atom 1st Ionization Energy kJ/mole F Cl Br I

A mole of atoms can be converted into a mole of +1 ions by applying the 1st ionization energy. Anything (almost) is possible. Making fluorine a +1 cation is possible, but hard. It requires a lot of energy compared to lithium. Noble gases can be forced into +1 cations too. That is really hard to do but possible. Atoms like Mg and Ca make +2 ions. Al makes a +3. To convert a mole of Mg into a mole of +2 ions requires the application of the 1st Ionization energy PLUS the application of the 2nd Ionization energy (the energy required to remove a second electron from a mole of +1 ions to make them form into +2 ions.) There is also a 3rd Ionization energy. (We concern ourselves only with the 1st level energy, not the others. I add it here because you are smart enough to understand it, and also smart enough to let it go). Noble gases tend to be “perfect” atoms, and it’s very hard to remove their electrons, as the Table S shows us with very high 1st ionization energies. The atom with the highest 1st ionization energy is helium, only two electrons but super small. It’s so hard to remove those electrons!

Metallic Property and Non-Metallic Property

Metals are on the left side of the Periodic Table of Elements. They have a variety of properties that make them “metallic”, such as: luster, electric conductivity, electric conductivity, malleable, ductile, low specific heat capacity, higher density, higher melting point, form cations, etc. Non-metals are on the right side of the table (plus Hydrogen). They have pretty much the opposite properties

• f metals. Nonmetal solids are brittle, not able to change shape. The gases are weirder. They also tend to

form anions, or no ions, they don’t conduct heat, don’t conduct electricity, and are dull rather than lustrous. If each property could somehow be ranked, if we “measure” all metals against each other, the MOST METALLIC METAL would be Francium, Fr. In fact, the closer to Fr a metal is on the table, the more metallic it is. Use this idea to rate or rank groups of

• metals. Example: polonium, lead, silver and zirconium are all metals. Zr is the “closest” to Fr on the table,

therefore, Zr is the most metallic of these four metals. Non-metals are on the right side of the table (plus Hydrogen). They have pretty much the opposite properties

• f metals. Nonmetal solids are brittle, not able to change shape. The gases are weirder. They also tend to

form anions, or no ions, they don’t conduct heat, don’t conduct electricity, and are dull rather than lustrous.

If these nonmetals were ranked (which is the most nonmetallic of them all?) Helium would come up as the MOST NON-METALLIC NON-METAL. The closer an atom is to He on the table, the more non-metallic it is said to be. Example, C, Cl and Ne are all non-metals and neon is the closest on the table to helium, so Ne is the most non-metallic of these three non-metals. Sometimes kids think of crazy questions, like which is more metallic, Cesium or radium, or which is more non-metallic, F or Ne. These questions cannot be answered by our simple distance evaluation, so don’t worry, no one will ever ask those questions of you.

Parts of the Periodic Table of Elements

Even the name of the table is important. The properties of the elements periodically repeat themselves, so that is why it’s the PERIODIC table. Group 1 = alkali metals Group 2 = alkaline Earth metals Group 17 = halogens Group 18 = Noble Gases hydrogen is the exception, it’s a non-metal that acts like a group 1 metal in bonding Transitional metals stretch the middle of the table, plus some under groups 13 to 16. Inner Transitionals are at the bottom of the chart, and all fit into GROUP 3, and periods 6 + 7 Under Y-39 fits 57-71. Under that comes 89-103. All of these are group 3 metals 7 Metalloids touch the staircase line from group 13 down into group 17, with 2 exceptions: Al & Po. All atomic masses are based against carbon-12, an atom with 6 protons and 6 neutrons. It’s said to have the exact mass of 12 amu. One AMU is one twelfth the mass of one C-12 atom. At STP, all metals are solids, except for Hg, which is liquid. At STP most non-metals are gases, but Br is a liquid and some are solids as well. The modern periodic table was first devised by Dimitri Mendeleev, a Russian chemist. It was a tremendous achievement of figuring out a pattern for 70 plus elements into a table that had no particular shape.

The Periodic Law When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their chemical and physical characteristics in the groups.

ALLOTROPES An allotrope is a pure form of an element, but it is bonded together in a different way, so it has different

• properties. Examples include CARBON: in the form of graphite, diamonds, and “Bucky balls”.

Another example is oxygen as O2 we breathe and O3 called ozone. Oxygen and ozone are allotropes of oxygen.

3 allotropes of carbon, All forms of the pure element, but in different structure, with different properties.

Oxygen atoms do not exist normally, they are too unstable. Both O2 + O3 (ozone) are allotropes of oxygen.

General Information about the Periodic Table

• 1. The Periodic Table was developed by _____________________________________, the Russian chemist.
• 2. It has ___________ groups and ______________ periods.

Name these parts of the periodic table

• 3. Group 1 is called the _______________________________________
• 4. Group 2 is called the ____________________________________________________
• 5. The center of the table are the _________________________________________________________
• 6. On the right side of the staircase are the _________________________________________________
• 7. Group 17 elements are called the _______________________________________
• 8. Group 18 elements are the ___________________________________________________
• 9. Seven of the nine atoms that touch the staircase are sort of “in between” metals and nonmetals, they are

called the ____________________________________.

• 10. List the metalloid atomic symbols _________________________________________________________
• 11. The “left over” atoms at the bottom of the table are called the

________________________________________________________________________ metals.

• 12. Get up now and go to the back of the room to look at that other periodic table. It shows much better what

atoms are in group 3, and how they should fit together. List the first 2 group 3 metal symbols: __________________

• 13. How many atoms are in Group 3—Period 5? ____________
• 14. How many atoms are in Group 3—Period 5? ____________
• 15. How many atoms are in Group 3 totally? ____________

• 16. The columns of the Periodic Table are the __________________________
• 17. The rows going across the Periodic Table are the __________________________
• 18. Atoms in the same group share many similarities of chemical properties, because they have similar

electron orbitals, which means they bond in similar ways. What do atoms in the same period have in common? _____________________________________________________________________________________

• 19. What does PERIODIC even mean here?
• 20. The elements of the Periodic Table are arranged in order of ___________________________
• 21. State the PERIODIC LAW
• 22. List ALL of the nonmetals (by symbol, in atomic number order) (there are 22 of them)
• 23. How many atoms are on the Periodic table? __________ How many are Metals? _____________.

Define these words…

• 24. FIRST IONIZATION ENERGY
• 25. ELECTRONEGATIVITY
• 26. ATOMIC RADIUS
• 27. NET NUCLEAR CHARGE
• 28. AVERAGE WEIGHTED ATOMIC MASS

——————— All atoms have isotopes, which are chemically identical atoms with a different number of neutrons. These are naturally occurring, and each element has two, up to a dozen different isotopes. Not all neon atoms are alike, they all are chemically identical, but they have different masses.

Isotope Mass number Atomic Mass Naturally Occurring Percentage ATOMIC MASS CALCUALTIONS Neon-20 20 19.9924 amu 90.480% = 18.088 Neon-21 21 20.9938 amu 0.270% = 0.0566 Neon-22 22 21.9914 amu 9.25% = 2.0342 100.0% Sum = 20.18 AMU Totals → → →→ →

Given the simple “rounded” atomic mass numbers and the actual scientific atomic masses, do not use the rounded numbers to calculate the average weighted atomic mass. Using these tables, calculate the average weighted atomic mass for these two atoms. ———————— Calculate the number of protons, neutrons, and electrons in the most common isotope

• f chromium. (mass of 51.996 rounds to 52 AMU, the most common isotope of

Cr has mass of 52. 52 = total number of protons plus neutrons. How many protons, neutrons & electrons in the most common isotope of lead. How many protons, neutrons, and electrons in the most common isotope of chlorine. Sometimes instead of the MOST COMMON isotope, you are told to Calculate the number of protons, neutrons, and electrons in a particular isotope, like, zinc-66. How many protons, neutrons, and electrons in the most common isotope of Zr-93. How many protons, neutrons, and electrons in the most common isotope of Ra-227. Isotope Mass number Atomic Mass Naturally Occurring Percentage ATOMIC MASS CALCUALTIONS Nitrogen-14 14 14.003074 99.63

=

Nitrogen-15 15 15.000108 1.11%

=

Totals → → →→ → 100.0%

Sum = Unknown Isotope “X” Mass number Atomic Mass Naturally Occurring Percentage ATOMIC MASS CALCUALTIONS X-126 126 125.979 amu 16.05% = X-127 127 126.994 amu 79.78% = X-128 128 127.967 amu 4.17% = 100.0% Sum = Totals → → →→ → Cr 54 Protons 24 Neutrons 54—24 = 30 Electrons 24 Pb-207 Chlorine-35 Zn-66 Zirconium-93 Ra-227 Number of protons Number of neutrons Number of electrons

Trends of the Periodic Table Homework # 1

• 1. Define net nuclear charge
• 2. What are the net nuclear charges for these 10 atoms? Mg ___ P ___ Sc ___ V ___

Au ___ Ir ____ Cr ____ Hg ____ Al ____ Pb ____

• 3. In a full sentence, state is the GROUP TREND for net nuclear charge.
• 4. In a full sentence, state the PERIOD TREND for the net nuclear charge.
• 5. Explain what happens with the ATOMIC MASS of Cobalt & Nickel. Does this destroy the general trend?
• 6. Using the boxes, figure out the group trend for atomic radius, then STATE it in one sentence.
• 7. Using the boxes, figure out the period trend for atomic radius, then STATE it in one sentence.
• 8. Explain why the group trend and the period trend for atomic radius occurs.

Group 2 atomic radius (pm) Be Mg Ca Sr Ba Period 3 Na Mg Al Si P S Cl atomic radius (pm) Ar

Trends of the Periodic Table Homework # 2

• 1. What is the most nonmetallic element on the periodic table? __________
• 2. What is the most metallic element on the periodic table? _________
• 3. Circle the most metallic of these three elements. Zinc Copper Iron
• 4. Circle the most non metallic of these three elements Aluminum Fluorine Sulfur
• 5. What is the technical name of the group 2 metals? _____________________________________________
• 6. What is the name of the group 17 nonmetals? ________________________________________________
• 7. What are the group 1 metals also known as? _________________________________________________
• 8. How many elements are in group 3? _________
• 9. List the symbols of ALL of the non metals: __________________________________________________

_____________________________________________________________________________________

• 10. List the symbols of the metalloids __________________________________________________________
• 11. What is a metalloid?
• 12. What element is in period 5, group 11? ______ What element is in period 4, group 16? ______
• 13. Groups 3-12 (and the “∆” of metals from Al to Tl to Po) make up what are known as the

___________________________________________________________ metals

• 14. How many protons, neutrons and electrons are in the element with the greatest density on the table?
• 15. What is the mass of the most common isotope of the element tantalum? ________________ amu

Trends of the Periodic Table Homework # 3

• 1. Define 1st Ionization Energy.
• 2. What this the GROUP TREND for 1st Ionization energy going down any group?
• 3. What is the PERIOD TREND for 1st Ionization energy going across any period?
• 4. What atoms have the highest 1st Ionization energy and WHY?
• 5. Define Electronegativity.
• 6. Define relative scale.
• 7. Define arbitrary scale.
• 8. Which element has the highest EN value? What does that mean about this atom?
• 9. Which part of the periodic table tends to have the highest EN values? Why?
• 10. Which part of the periodic table tends to have very low EN values? Why?

Trends of the Periodic Table Homework # 4

• 1. Which is bigger or smaller, the Na atom or the Na+1 cation? Say it in a sentence.
• 2. Which is bigger or smaller, the Mg atom or the Mg+2 cation? Say it in a sentence.
• 3. Which is bigger or smaller, the Al atom or the Al+3 cation? Say it in a sentence.
• 4. Which is bigger or smaller, the N atom or the N-3 anion? Say it in a sentence.
• 5. Which is bigger or smaller, the O atom or the O-2 anion? Say it in a sentence.
• 6. Which is bigger or smaller, the F atom or the F-1 anion? Say it in a sentence.
• 7. Why are cations smaller than their atoms, and anions larger than their atoms - all of the time?
• 8. State the GROUP TREND for cation size. Full sentence only for credit.
• 9. State the PERIOD TREND for cation size. Full sentence only for credit.
• 10. State the GROUP TREND for anion size. Full sentence only for credit.
• 11. State the PERIOD TREND for anion size. Full sentence only for credit.

Trends of the Periodic Table Homework # 5 This bar graph shows the number of TV’s produced in Japan by specific years.

• 1. State the trend of tv production over time in Japan.
• 2. During 2016-2018 the number of tv’s produced stays almost flat. What does that do to the trend?

Table G, from your reference table, shows the number of grams

• f ten different solutes dissolving into water, at all temperatures

from zero to 100C. Answer these simple questions.

• 3. What is the trend of grams of ammonium chloride

dissolving into water that is getting hotter.

• 4. What is the trend of grams of ammonia dissolving into

water that is getting hotter.

• 5. What is the trend of grams of potassium nitrate dissolving

into water that is cooling down? (wake up here) 10 9 8 7 6 5 4 3 2 1 2013 2014 2015 2016 2017 2018 2019 2020 100,000’s of flat screen TV’s produced in Japan over 7 years (2020 is expected production)

When it comes to the Trends of the Periodic Table , I can…

• 1. I can classify elements as metals, nonmetals, or metalloids based on their placement on

the Periodic Table. Classify each of the following elements as a metal (M), nonmetal (NM), or metalloid (LOID).

• 2. I can state the group names for groups 1, 2, 17, and 18.
• 3. I can explain why elements in the same group have similar chemical properties.
• 4. I can explain why the elements in Group 18 don’t usually react with the other elements.
• 5. I can state the meaning of “STP” and where on the Reference Table it can be found.
• 6. I can state the names and symbols for the 2 elements on the Periodic Table that are liquids at STP.
• 7. I can state the names and symbols of the 11 elements that are always gases at STP.
• 8. I can state how the elements on the Periodic Table are arranged.
• 9. I can list the 7 diatomic elements.
• 10. I can define electronegativity, arbitrary + relative scales, 1st ionization energy, atomic radius, ionic radius,

net nuclear charge, metallic + non-metallic character, and allotrope.

• 11. I can state the group trend + period trend for electronegativity and explain why it occurs.
• 12. I can state the group trend + period trend for 1st ionization energy + explain why it occurs.
• 13. Skip this one here.
• 14. I can state the group trend and period trend for atomic radius and explain why it occurs.
• 15. I can state the group trend + period trend for metallic character + explain why it occurs.
• 16. I can state the group trend and the period trend for non metallic character, and I can explain why it occurs.
• 17. I can state the group trend and period trend for net nuclear charge and why it occurs.
• 17. I can state the group trend and period trend for cation radius and why it occurs
• 18. I can state the group trend and period trend for anion radius and why it occurs
• 19. I can list 10 properties of metals.
• 20. I can list 8 properties of nonmetals.

Answers are all on Arbuiso.com on the Periodic Table page.

B K Li C Ar Sb H Fe Au S F Si Fr He Rn Ge Al As Bi I