[PDF] Element Entry Identification What is the Atomic Number? Terms

141038_7PeriodicTablePart2Handout.pdf 1 Name_______________________________ Period:________ Using the terms answer the following questions regarding global winds 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. Horizontal rows are called _______________ Numbered ____ Vertical columns are called _____________ Numbered _____ ment Entry Identificat

What is the Atomic Number?

Equator Polar Easterlies

Westerlies Tradewinds

Doldrums Heated area

Coriolis Effect up

Air Pressure (high, low)

Earth"s rotation

2 . 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 neutrons and thus different weights (Mass number = A). The mass number is the total number of protons and neutron in the nucleus. Atoms are said to be isotopes i f 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 carbon atoms they have the same number of protons: 6. (The atomic number of carbon is

6.) All elements of carbon have six 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 isotope can be tracked.

Another way of showing isotopes is ܺ

௓஺ where X is some element"s symbol, A is the elements mass number and Z is the elements atomic number. How many protons _________ What is the atomic number _____

What is the mass number _____

What is the number of neutrons _____

There are three isotopes of hydrogen that all differ by the number of neutrons since all hydrogens have the same number of protons.

NORMAL HYDROGEN H

-l Most hydrogen atoms consist of just a single proton and an electron... 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 deuterium. These atoms are twice as heavy as “normal" hydrogen atoms because they contain a neutron in addition to the proton in normal H. TRITIUM H-3 These atoms contain a proton and two neutrons. Draw diagrams showing normal hydrogen, deuterium and tritium label protons (+), neutrons (0) and electrons (-) indicating their respective charges.

Normal Hydrogen

Deuterium Tritium 3 Complete the Isotope Table to identify the number of subatomic particles uranium-235 uranium-238 ܫ ଵଷଵ ܫ ହଷ ଵଶ଻

Figure out the element from subatomic

clues and practice writing isotopic notations ܺ ௓ ஺

Cadmium

-116 49 113
45 103
197

Au

Xenon-136

40 71
106 180
33 42

Mercury-204

3

H

Calculation Space

4

Relative Atomic Mass

Describe the

masses of protons and neutrons in comparison to electrons ___________________________________________________________________________________ ___________________________________________________________________________________ ___________________________________________________________________________________ ___________________________________________________________________________________

Reactivity relates to electrons

Name of

Group

Lewis Dot Structure of

Valence Electrons

alkali metals alkaline earth metals

Boron group

Carbon group

Nitrogen

Group

Oxygen Group

Chacagon

halogens noble gases Atoms, tend to take on the lowest-energy, most stable configuration they can. This is why the

electron shells of an atom 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.

In general, atoms are most stable,

least reactive, when their outermost electron shell 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 octet rule. 5 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. 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

Simply take a look at the period number - it is the number of the shell in given atom. Lithium can be

found in the second period so electrons will be placed on two shells Finally Bohr"s atomic model of lithium will look like this:

Figure 1

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): ܥ ଵଶ ଺ ܤ ସ ଽ ܤ ହ ଵ଴ Aluminum -27

Second orbital n=2

Third orbital n=3

6 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.

Sodium (Na)__________

Potassium (K)________ Helium (He)_______ Carbon (C) ________ Nitrogen (N) _______ Oxygen (O) _____ Chlorine (Cl)_________ Phosphorus (P)____________ Sulfur (S)_________

7 Summary

1. How many electrons can each

level hold ? 1 st _____ 2 nd ______ 3 rd ______

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

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? 8 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. 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.

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 ion, 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 electron configuration 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. The filling of electrons in a dot diagram is accomplished by putting one dot on each of four sides before doubling up. bond. These values are based on an arbitrary scale. Electronegativity can also be described as electroaffinity.

11. Bonding guidelines:

Metals react with nonmetals 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 9

Bonding

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15

2. Electron ___

____ __ model - Proposes that all metal 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

reacta nts 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.

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. 10

Ionic Bonds

Covalent Bonds

Metallic Bonds

11 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. Ionization energy - energy required to remove an electron from an ion or gaseous atom Atomic radius - 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 Electron affinity - ability of an atom to accept an electron

The periodicity of th

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

ĺĺ

Ionization Energy Increases Ionization Energy Decreases Electronegativity Increases Electronegativity Decreases Atomic Radius Decreases Atomic Radius Increases 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 leve

l.

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 12

Ionic radius

Quick Write. Explain the difference between an ion and an atom of the same element ___________________________________________________________________________ _______________ __________________________________________________________ _____________ ____ ______________ 4. a. b. c. d. +1 5.

The octet rule

octet rule. gain or lose how many electrons and the resultant charge a. b. c. d. 7. a. b. c. d. 13

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 fluorine, oxygen, and chlorine, 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 sodium, potassium, and strontium, 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. For each of the following pairs, predict which atom has the higher first ionization energy.

Mg, Na S, O Ca, Ba Cl, I

For each of the following pairs, predict which atom forms a positive ion more easily.

Be, Ca F, I Na, Si K, Ca

Explain how you can predict which atoms will form cations more easily than anions __________________________________________________________________________________________ ___________________________________________________ ________________ _______________________ 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, rubidium, and barium, 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. Electronegativities

generally increa se across a period and decrease down through a group or increase ______________________________________. For each of the following pairs, predict which atom has the higher electronegativity.

Mg, Na Ca, Ba

Na, Al S, O

Cl, I Se, Br

14 15 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 neon, 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 nucle us. Atomic sizes generally decrease from left to right and increase from top to bottom for the same reasons.

1. On the top, plot a graph of ionization energy (y-axis

left) vs. atomic number (x-axis) and middle plot electronegativity (y axis right) and atomic number (x- axis). On the bottom plot a separate graph of atomic radius vs. atomic number. For each graph connect successive dots with straight lines. Also, ensure that identical atomic numbers a re plotted on the same vertical position on the sheet (i.e. atomic number 1 in the top graph should be on the same line as atomic number 1 in the bottom graph). a) Examine your graph of ionization energy (IE) vs. atomic number. Which elements are found at the main peaks on your graph (there should be 3) and what do these elements have in common? b) Which elements are found at the main valleys on IE vs atomic number graph (there should be 3) and what do these elements have in common? c) Examine your graph of atomic radius verses atomic number. Which elements are found at the peaks on your 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)?

Atomic

Number Element Symbol Ionization Energy, 1st

(kJ/mol) Electro- negativity Atomic

Radius

(pm)

1 H 1312 2.2 32

2 He 2372 31

3 Li 520 0.98 123

4 Be 899 1.57 90

5 B 801 2.04 82

6 C 1086 2.55 77

7 N 1402 3.04 75

8 O 1314 3.44 73

9 F 1681 4.00 72

10 Ne 2081 71

11 Na 496 0.93 154

12 Mg 738 1.31 136

13 Al 578 1.61 118

14 Si 786 1.90 111

15 P 1012 2.19 106

16 S 1000 2.58 102

17 Cl 1251 3.16 99

18 Ar 1521 98

19 K 419 0.82 203

20 Ca 590 1.00 174

Periodic Table of the Elements (Used for Grade 8 and High School) California Science Test Copyright © 2018 California Department of Education30 zinc 65.38
31
gallium 69.72
32
germanium 72.63
33
arsenic 74.92
21
scandium 44.96
20 calcium 40.08

Atomic NumberAtomic Number

Element Name

* If this number is in parentheses, then it refers to the atomic mass of t he most stable isotope.

Element Symbol

Average Atomic Mass

1 hydrogen 1.01 1 hydrogen 1.01 3 lithium 6.94 11 sodium 22.99
19 potassium 39.10
37
rubidium 85.47
55
cesium

132.91

87
francium (223) 4 beryllium 9.01 12 magnesium 24.30
38
strontium 87.62
56
barium

137.33

88
radium (226) 39
yttrium 88.91
57
lanthanum

138.91

63
europium

151.96

89
actinium (227) 95
americium (243) 57-71

89-10322

titanium 47.87
40
zirconium 91.22
72
hafnium

178.49

104
rutherfordium (267) 58
cerium

140.12

64
gadolinium

157.25

90
thorium

232.04

96
curium (247) 23
vanadium 50.94
41
niobium 92.91
73
tantalum

180.95

105
dubnium (268) 59
praseodymium

140.91

65
terbium

158.93

91
protactinium

231.04

97
berkelium (247) 24
chromium 52.00
42
molybdenum 95.95
74
tungsten

183.84

106
seaborgium (269) 60
neodymium

144.24

66
dysprosium

162.5069

thulium

168.93

92
uranium

238.03

98
californium (251) 101
mendelevium (258) 25
manganese 54.94
43
technetium (98) 75
rhenium

186.21

107
bohrium (270) 61
promethium (145) 67
holmium

164.93

70
ytterbium

173.05

93
neptunium (237) 99
einsteinium (252) 102
nobelium (259) 26
iron 55.85
44
ruthenium

101.07

76
osmium

190.23

108
hassium (269) 62
samarium

150.36

8 erbium

167.26

71
lutetium

174.97

94
plutonium (244) 100
fermium (257) 103
lawrencium (266) 27
cobalt 58.93
45
rhodium

102.91

77
iridium

192.22

109
meitnerium (278) 28
nickel 58.69
46
palladium

106.42

78
platinum

195.08

110
darmstadtium (281) 29
copper 63.55
47
silver

107.87

79
gold

196.97

111
roentgenium (282) 48
cadmium

112.4180

mercury

200.59

112
copernicium (285) 5 boron 10.81 13 aluminum 26.98
14 silicon 28.09
15 phosphorus 30.97
16 sulfur 32.06
17 chlorine 35.45
18 argon 39.95
6 carbon 12.01 7 nitrogen 14.01 8 oxygen 16.00 9 19.00 10 neon 20.18
2 helium 4.00 49
indium

114.82

81
thallium

204.38

113
nihonium (286) 50
en-US tin

118.71

82
lead

207.21

114
(289) 51
antimony

121.76

83
bismuth

208.98

115
moscovium (289) 34
selenium 78.97
52
tellurium

127.60

84
polonium (209) 116
livermorium (293) 35
bromine 79.90
53
iodine

126.90

85
astatine (210) 117
tennessine (294) 36
krypton 83.80
54
xenon

131.29

86
radon (222) 118
oganesson (294)