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Periodic Table Part 2 Name______Period:______4

Learning Objective

Using the terms answer the following questions regarding global winds

Terms Equator Polar Easterlies

Westerlies Tradewinds

Doldrums Heated area

Coriolis Effect up

Air Pressure (high, low)

Earth’s rotation

 Air currents between 0 and 30 latitude are called ______ Air current that brings the weather to North America and Canada between 30 and 60 degrees latitude is called ______.  Air currents between 60 and 90 degrees latitude are called ______ Right near the equator is a windless area called the ______ Zero degrees latitude is called the ______ What factor has the greatest effect on wind speed______. Winds move from ______to ______.  The deflection of winds in the northern hemisphere to the right is called the ______and is due to Earth’s ______ ______rises and can be pushed up and forced ______by cold air.

Periodic Table Horizontal rows are called ______Numbered ____

Vertical columns are called ______Numbered _____

How is the periodic table Element Entry Identification organized?______. Then after the primary organization the periodic table is organized by ______What is the ?

1

ISOTOPES All of any given element have the same numbers of (atomic number = Z) in their nucleus. Atoms are identified based on the number protons in the nucleus. The Periodic Table is organized in order of increasing atomic number and chemical reactivities.

However, atoms of the same element may have different numbers of and thus different weights ( = A). The mass number is the total number of protons and in the nucleus. Atoms are said to be if they are of the same element but they have different masses (weights) due to different numbers of neutrons. C-l2 and C-14 are isotopes. Since both are elemental atoms they have the same number of protons: 6. (The atomic number of carbon is 6.) All elements of carbon have six protons.

Atoms of C-l2, like any carbon atoms must have 6 protons. In order for these atoms to have a mass number of 12 they must also contain 6 neutrons (6 protons + 6 neutrons). Atoms of C-14 must also have 6 protons (all carbon atoms do). However, in order for these atoms to have a mass of 14 they must contain _____ protons + ______neutrons.

Generally, isotopes of an element behave identically in terms of how they react with other chemicals. The only difference is in their weights. Isotopes are not present in the same proportion in nature, that is some isotopes are more common than others. Their difference in weight makes they very useful when doing things like medical tests looking how the body responds because an uncommon can be tracked. ISOTOPIC NOTATION Another way of showing isotopes is where X is some element’s symbol, A is the elements mass number𝐴𝐴 and Z is the elements atomic number. 𝑍𝑍𝑋𝑋 Mass Number = Protons +Neutrons Example How many protons ______What is the atomic number _____ What is the mass number _____ What is the number of neutrons _____

Hydrogen Isotopes There are three isotopes of that all differ by the number of neutrons since all have the same number of protons. NORMAL HYDROGEN H-l Most hydrogen atoms consist of just a single and an ... no neutrons; thus H-1 has a mass of 1amu or a mole of H-1 has a mass of 1 gram. About 99.98% of all hydrogen atoms are normal hydrogen; (sometimes called protium). HEAVY HYDROGEN H-2 Sometimes called . These atoms are twice as heavy as “normal” hydrogen atoms because they contain a neutron in addition to the proton in normal H. H-3 These atoms contain a proton and two neutrons. Draw diagrams showing normal hydrogen, deuterium and tritium label protons (+), neutrons (0) and (-) indicating their respective charges. Normal Hydrogen Deuterium Tritium

2 Complete the Isotope Table to identify the number of subatomic particles # of # of # of Isotope name # # protons neutrons electrons

-235

uranium-238

131 𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼 127 53𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼

Figure out the element from subatomic clues and practice writing isotopic notations

ISOTOPE NUMBER NUMBER MASS Average

OF OF NUMBER Atomic PROTONS NEUTRONS Mass 𝐴𝐴 𝑍𝑍𝑋𝑋 -116

49 113

45 103

197Au

Xenon-136

40 71

106 180

33 42

Mercury-204

3H

Calculation Space

3

Relative Atomic Mass Describe the masses of protons and neutrons in comparison to electrons

______Reactivity relates to electrons Representative Elements Name of Group Example Number of Valence Lewis Dot Structure of Group # (symbol) Electrons Valence Electrons

alkali

alkaline earth metals

Boron group

Carbon group

Nitrogen

Group

Oxygen Group

Chacagon

noble gases

Atoms, tend to take on the lowest-energy, most stable configuration they can. This is why the electron shells of an are filled from the inside out, with electrons filling up the low-energy shells closer to the nucleus before they move into the higher-energy shells further out. The number of electrons in the outermost shell of a particular atom determines its reactivity, or tendency to form chemical bonds with other atoms. This outermost shell is known as the valence shell, and the electrons found in it are called valence electrons. In general, atoms are most stable, least reactive, when their outermost is full. Many of the elements need eight electrons in their outermost shell in order to be stable, and this rule of thumb is known as the .

4 The structure of the atom - How to draw the Bohr atomic models Atoms are composed of three basic subatomic particles: protons, neutrons, and electrons. In 1913, Niels Bohr came up with a new atomic model. Protons with positive charge and neutrons with neutral charge are located in the nucleus of the atom in the center. Strong nuclear forces hold them together. Electrons orbit around the nucleus on fixed distances called shells. Electrons are negative First shell (the closest to the nucleus) can hold up to two electrons, second eight electrons, third 18, fourth 32.

The last 2 things to remember is that electrons will fill the closest to the nucleus shells first and electrons are much smaller compared to protons and neutrons.

How to determine the number of protons, neutrons, and electrons in neutral atom • The number of protons in the nucleus of the atom is equal to the atomic number (Z). • The number of electrons in a neutral atom is equal to the number of protons. • The mass number (A) of the atom is equal to the sum of the number of protons and neutrons in the nucleus. To find number of neutrons simply subtract number of protons from the mass number

How to find the number of shells? Simply take a look at the number - it is the number of the shell in given atom. can be found in the second period so electrons will be placed on two shells

Bohr’s atomic model Finally Bohr’s atomic model of lithium will look like this: Figure 1

Orbital How many total electrons First closest to nucleus

n=1 Second orbital n=2

Third orbital n=3

Exercise Now it’s your turn: Try to find the number of protons, neutrons and electrons and draw the Bohr’s atomic models for given elements (use periodic table):

Aluminum -27 6 9 10 12𝐶𝐶 4𝐵𝐵𝐵𝐵 5𝐵𝐵

5

For each element, write the total number of electrons on the line. Then color the correct number of electrons for each orbit. Remember, fill the orbit closest to the nucleus first, but never exceed the number each orbit can hold. Check the Periodic Table to find out how many electrons each element actually has.

Sodium (Na)______Potassium (K)______Helium (He)______

Carbon (C) ______Nitrogen (N) ______Oxygen (O) _____

Chlorine (Cl)______Phosphorus (P)______Sulfur (S)______

6 Summary

1. How many electrons can each level hold ? 1st _____ 2nd______3rd ______

2. What term is used for the lelectrons in the outermost shell or ender level?______

3. Scientists use two types of diagrams to show the electron confiruration for atoms. Using the Periodic Table complete the following diagrams showing a Bohr Model and a Lewis Structure

Bohr Model Lewis Dot Structure All Electrons (valence electrons) Atomic Number = 16 Mass Number = 32 Protons =

Neutrons =

Electrons =

Draw the Bohr Model and Lewis Dot structures for the following elements

Which elements had a filled outermost shell______

Which element would be most likely to lose electrons in a chemical bond______

Which element would be most likely to gain electrons in a chemical bond

Which elements are not likely to bond with other elements?

7 1. Chemical compounds are formed when atoms are bonded together  Breaking a chemical bond is an endothermic process.  Forming a chemical bond is an exothermic process.  Compounds have less potential energy than the individual atoms they are formed from. 2. Two major categories of compounds are ionic and molecular (covalent) compounds. 3. Compounds can be differentiated by their chemical and physical properties.  Ionic substances have high melting and boiling points, form crystals, dissolve in water (dissociation), and conduct electricity in solution (aqueous) and as a liquid, not as a solid.  Covalent or molecular substances have lower melting and boiling points, do not conduct electricity.  Polar substances are dissolved only by another polar substance (water soluble). Non-polar substances are dissolved only by other non-polar substances (oil soluble).  Metallic compounds have high melting and boiling points and conduct heat and electricity as a solid or liquid or aqueous. 4. Chemical bonds are formed when valence electrons are:  Transferred from one atom to another – ionic.  Shared between atoms – covalent.  Mobile in a free moving “sea” of electrons – metallic. 5. In multiple (double or triple) covalent bonds more than 1 pair of electrons are shared between two atoms. 6. When an atom gains an electron, it becomes a negative , anion, and its radius increases. 7. When an atom loses an electron, it becomes a positive ion, cation, and its radius decreases. 8. Atoms gain a stable by bonding with other atoms.  Atoms are stable when they have a full valence level.  Most atoms need 8 electrons to fill their valence level.  H and He only need 2 electrons to fill their valence level.  The noble gasses (group 18) have filled valence levels. They do not normally bond with other atoms. 9. Electron-dot diagrams (Lewis structures) represent the valence electrons in elements, compounds and .  The filling of electrons in a dot diagram is accomplished by putting one dot on each of four sides before doubling up.  10. indicates how strongly an atom of an element attracts electrons in a chemical bond. These values are based on an arbitrary scale. Electronegativity can also be described as electroaffinity. 11. Bonding guidelines:  Metals react with to form ionic bonds, ionic compounds, formula units (salts).  Nonmetals bond with nonmetals to form covalent compounds (molecules).  Ionic compounds with polyatomic ions have both ionic and covalent bonds.  Metals react with Metals to form metallic bonds

15. Physical properties of a substance can be explained in terms of chemical bonds and intermolecular forces. These include conductivity, malleability, solubility, ductility, hardness, and . 8 Bonding 1 2 3

4 5

6 7 8 9

10 11

12 13

14 15

Across 2. Electron ______model - Proposes that all atoms in a metallic solid contribute their valence electrons to form a "sea" of electrons; can explain properties of metallic solids. 7. A chemical reaction in which a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the product molecules. 9. ______bond - The force that holds two atoms together; may form by the attraction of a positive ion for a negative ion or by the attraction of a positive nucleus for negative electrons (electrostatic). 10. ______energy - The energy required to separate one mole of the ions of an ionic compound, which is directly related to the size of the ions bonded and is also affected by the charge of the ions. 11. An ion that has a positive charge; forms when valence electrons are removed, giving the ion a stable electron configuration. 12. ______unit - The simplest ratio of ions represented in an ionic compound. 13. ______covalent - A type of bond that forms when electrons are not shared equally. 14. An ion that has a negative charge; forms when valence electrons are added to the outer energy level, giving the ion a stable electron configuration. 15. ______structure - Uses an electron-dot diagram to show how electrons are arranged in molecules.

Down 1. ______bond - The attraction of a metallic cation for delocalized electrons. 3. ______bond - A chemical bond that results from the sharing of valence electrons. 4. Ionic ______- The electrostatic force that holds oppositely charged particles together in an ionic compound. 5. A chemical reaction in which more energy is released than is required to break bonds in the initial reaction. 6. ______electrons - The electrons involved in metallic bonding that are free to move easily from one atom to the next throughout the metal and are not attached to a particular atom. 8. ______ion - An ion formed from only one atom.

9 Bonding

Ionic Bonds

Covalent Bonds

Metallic Bonds

10 TRENDS in the Periodic Table Periodicity refers to the recurring trends that are seen in the element properties.

The electron structure of an atom determines many of its chemical and physical properties. Because the periodic table reflects the electron configurations of the elements, the table also reveals trends in the elements’ chemical and physical properties.

What Are the Periodic Properties?

– energy required to remove an electron from an ion or gaseous atom  – half the distance between the centers of two atoms that are touching each other  Electronegativity – measure of the ability of an atom to form a chemical bond  – ability of an atom to accept an electron

The periodicity of these properties follows trends as you move across a row or period of the periodic table or down a column or group:

Moving Left → Right Moving Top → Bottom Ionization Energy Increases Ionization Energy Decreases Electronegativity Increases Electronegativity Decreases Atomic Radius Decreases Atomic Radius Increases

Atomic radius The atomic radius is a measure of the size of an atom. The larger the radius, the larger is the atom. Research shows that atoms tend to decrease in size across a period because the nuclei are increasing in positive charge and the increased nuclear charge pulls the outermost electrons closer to the nucleus, making the atom smaller. Moving down through a group, atomic radii increase. Even though the positive charge of the nucleus increases, each successive element has electrons in the next higher energy level. Electrons in these higher energy levels are located farther from the nucleus than those in lower energy levels. The increased size of higher energy level outweighs the increased nuclear charge. Therefore, the atoms increase in size as the energy level (period) increases.

1. For each of the following pairs, circle which atom is larger.

a. Mg, Sr c. Ge, Sn e. Cr, W

b. Sr, Sn d. Ge, Br

11 2. Comparing elements from left to right across a period, what general trend would you predict for the atomic radius? Explain the basis for your prediction.

______3. Comparing elements from down to up a group, what general trend would you predict for the atomic radius? Explain the basis for your prediction. ______Ionic radius When an atom gains or loses one or more electrons, it becomes an ion. Because an electron has a negative charge, gaining electrons produces a negatively charged ion, whereas losing electrons produces a positively charged ion. As you might expect, the loss of electrons produces a positive ion with a radius that is smaller than that of the parent atom. Conversely, when an atom gains electrons, the resulting negative ion is larger than the parent atom. Practically all of the elements to the left of group 4A of the periodic table commonly form positive ions. As with neutral atoms, positive ions become smaller moving across a period and become larger moving down through a group. Most elements to the right of group 4A (with the exception of the noble gases in group 8A) form negative ions which are considerably larger than the positive ions to the left, also decrease in size moving across a period. Like the positive ions, the negative ions increase in size moving down through a group. When comparing negative ions in the same period, the more negative the larger the ion. Quick Write. Explain the difference between an ion and an atom of the same element ______

4. For each of the following pairs, predict which atom or ion is larger. a. Mg, Mg+2 b. S, S-2 c. Ca+2, Ba+2 d. Na+1, Al+3

5. Predict which of the ions, Mg+2 or S-2 is larger. Explain your prediction. ______

The octet rule When atoms lose or gain electrons, they generally do so until the ion has eight valence electrons—the stable electron configuration of a . This principle is called the octet rule. Exceptions to this rule are hydrogen, which can gain an electron, obtaining the stable configuration of , and elements in period 2, such as lithium and that lose electrons, also obtaining the helium configuration. The octet rule lets you predict the ionic charge of a representative element. For example, you can predict that an element in group 6A, having a high ionization energy, will gain two electrons to achieve a stable octet .

6. For each of the following elements, state whether it is more likely to gain or lose electrons to form a stable octet configuration and how many electrons will be gained or lost and the resultant charge. a. K b. Br c. O d. Mg

7. Which noble-gas configuration is each of the following elements most likely to attain by gaining or losing electrons? a. K b. Br c. O d. Be

12 Ionization energy Energy is required to pull an electron away from an atom. The first ionization energy of an element is the amount of energy required to pull the first valence electron away from an atom of the element. Atoms with high ionization energies, such as , , and , are found on the right side of the periodic table and are unlikely to form positive ions by losing electrons. Instead, they usually gain electrons, forming negative ions.

Atoms with low ionization energies, such as , , and , lose electrons easily to form positive ions and are on the left side of the periodic table. Recall that atoms decrease in size from left to right across a period. First ionization energies generally increase across a period of elements primarily because the electrons to be removed are successively closer to the nucleus. First ionization energies decrease moving down through a group of elements because the sizes of the atoms increase and the electrons to be removed are farther from the nucleus.

8. For each of the following pairs, predict which atom has the higher first ionization energy. a. Mg, Na b. S, O c. Ca, Ba d. Cl, I

7. For each of the following pairs, predict which atom forms a positive ion more easily. a. Be, Ca b. F, I c. Na, Si d. K, Ca

Explain how you can predict which atoms will form cations more easily than anions

______

Electronegativity When atoms combine chemically with each other, they do so by forming a chemical bond. This bond involves either the transfer of electrons or sharing of electrons to varying degrees. The nature of the bond between two atoms depends on the relative ability of each atom to attract electrons from the other, a property known as electronegativity.

The maximum electronegativity value is 3.98 for fluorine, the element that attracts electrons most strongly in a chemical bond. The trends in electronegativity in the periodic table are generally similar to the trends in ionization energy. The lowest electronegativity values occur among the elements in the lower left of the periodic table. These atoms, such as cesium, , and , are large and have few valence electrons, which they lose easily. Therefore, they have little attraction for electrons when forming a bond.

Elements with the highest electronegativity values, such as fluorine, chlorine, and oxygen, are found in the upper right of the periodic table (excluding, of course, the noble gases, which do not normally form chemical bonds). These atoms are small and can gain only one or two electrons to have a stable noble-gas configuration. Therefore, when these elements form a chemical bond, their attraction for electrons is large. generally increase across a period and decrease down through a group or increase ______.

9. For each of the following pairs, predict which atom has the higher electronegativity. a. Mg, Na d. Ca, Ba b. Na, Al e. S, O c. Cl, I f. Se, Br 13

14 Trends in the Periodic Table The ionization energy or ionization potential is the energy necessary to remove an electron from the neutral atom. It is a minimum for the alkali metals which have a single electron outside a closed shell. The first ionization energy, the energy to remove the first electron, generally increases across a row on the periodic maximum for the noble gases which have closed shells. An element often has multiple ionization energies, which correspond to the energy needed to remove first, second, third, and so forth electrons from the atom. For example, sodium requires only 496 kJ/mol to ionize it while , the noble gas immediately preceding it in the periodic table, requires 2081 kJ/mol. The ionization energy can be thought of as a kind of counter (opposite) property to electronegativity. A low ionization energy implies that an element readily gives electrons to a reaction, while a high electronegativity implies that an element strongly seeks to take electrons in a reaction. Electronegativity is a measure of the ability of an atom of an element to attract electrons toward itself in a chemical bond. The values of electronegativity calculated for various elements range from one or less for the alkali metals to three and one-half for oxygen to about four for fluorine. Ionization energy is the energy it takes to remove an electron from an atom. Generally in the periodic table, ionization energy and electronegativity increase from left to right because of increasing numbers of protons and decrease from top to bottom owing to an increasing distance between electrons and the nucleus. The atomic radius increases because the electron cloud which makes up most of the size gets bigger as the electrons fill energy levels away from the nucleus. Atomic sizes generally decrease from left to right and increase from top to bottom Atomic Element Ionization Electro- Atomic for the same reasons. Number Symbol Energy, 1st negativity Radius (kJ/mol) (pm) 1. On the top, plot a graph of ionization energy (y-axis 1 H 1312 2.2 32 left) vs. atomic number (x-axis) and middle plot 2 He 2372 31 electronegativity (y axis right) and atomic number (x- 3 Li 520 0.98 123 axis). On the bottom plot a separate graph of atomic 4 Be 899 1.57 90 radius vs. atomic number. For each graph connect 5 B 801 2.04 82 successive dots with straight lines. Also, ensure that 6 C 1086 2.55 77 identical atomic numbers are plotted on the same 7 N 1402 3.04 75 vertical position on the sheet (i.e. atomic number 1 in the 8 O 1314 3.44 73 top graph should be on the same line as atomic number 1 in the 9 F 1681 4.00 72 bottom graph). 10 Ne 2081 71 a) Examine your graph of ionization energy (IE) vs. 11 Na 496 0.93 154 atomic number. Which elements are found at the main 12 Mg 738 1.31 136 peaks on your graph (there should be 3) and what do 13 Al 578 1.61 118 these elements have in common? 14 Si 786 1.90 111

15 P 1012 2.19 106 b) Which elements are found at the main valleys on IE vs 16 S 1000 2.58 102 atomic number graph (there should be 3) and what do 17 Cl 1251 3.16 99 these elements have in common? 18 Ar 1521 98 c) Examine your graph of atomic radius verses atomic 19 K 419 0.82 203 number. Which elements are found at the peaks on your 20 Ca 590 1.00 174 graph and what do these elements have in common? d) Which elements are found at the valleys on your graph? What do these elements have in common?

e) How are atomic radii and ionization energy related (i.e. as atomic radius increases, what happens to the ionization energy)?

15 Periodic Table of the Elements (Used for Grade 8 and High School) Science Test

1 18 1 Legend 2 1 H He hydrogen Atomic Number 1 helium 1.01 13 14 15 16 17 4.00 2 H Element Symbol 3 4 Element Name hydrogen 5 6 7 8 9 10 Li Be 1.01 Average Atomic Mass B C N O F Ne 2 lithium beryllium carbon oxygen fluorine neon 6.94 9.01 * If this number is in parentheses, then it refers to the atomic mass of the 10.81 12.01 14.01 16.00 19.00 20.18 11 12 most stable isotope. 13 14 15 16 17 18 Na Mg Al Si P S Cl Ar 3 sodium aluminum sulfur chlorine 22.99 24.30 3 4 5 6 7 8 9 10 11 12 26.98 28.09 30.97 32.06 35.45 39.95 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 4 potassium krypton 39.10 40.08 44.96 47.87 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.38 69.72 72.63 74.92 78.97 79.90 83.80

37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 5 rubidium strontium cadmium xenon 85.47 87.62 88.91 91.22 92.91 95.95 (98) 101.07 102.91 106.42 107.87 112.41 114.82 118.71 121.76 127.60 126.90 131.29

55 56 57–71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 6 cesium barium 132.91 137.33 178.49 180.95 183.84 186.21 190.23 192.22 195.08 196.97 200.59 204.38 207.21 208.98 (209) (210) (222)

87 88 89–103 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og 7 (223) (226) (267) (268) (269) (270) (269) (278) (281) (282) (285) (286) (289) (289) (293) (294) (294)

57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 138.91 140.12 140.91 144.24 (145) 150.36 151.96 157.25 158.93 162.50 164.93 167.26 168.93 173.05 174.97

89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr uranium (227) 232.04 231.04 238.03 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (266)

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