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Unit 2: Atoms, Bonds,

Ions, and Their Properties

Name:____________________________________ Date:________ Period:_______

SECTION 1: ENGAGE

Investigation 1.1: What is an Atom?

SECTION 1: EXPLORE

Investigation 2.1: Introduction to the Periodic Table

SECTION 3: EXPLAIN

Investigation 3.1: Atomic Puzzle

Investigation 3.2: Understanding an Electron’s Behavior—Photons

Investigation 3.3: Electron Configuration

Investigation 3.4: Describing Valence Electrons using Lewis Dot Structures

Investigation 3.5: Understanding the formation of bonds and ions

SECTION 4: ELABORATE

Investigation 4.1: Replacement Reaction

Investigation 5.1: Ionic Bonds

Investigation 6.1: Covalent Bonds

Investigation 7.1:

SECTION 5: EVALUATE

Investigation 6.1: Test

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Investigation 1.1 What is an Atom?

Atoms are the building blocks of matter. To begin our investigation, you will be

defining and identifying all the parts of an atom.

1. Hypothesize what you think the following terms mean:

Nucleus:

Protons:

Electrons:

Neutrons:

2. In the space below, draw what you think an atom looks like. Please include

labels for each of the above terms.

3

Background Information:

All matter is made up of atoms. The center of the atom is the nucleus which is a

cluster of protons and neutrons. The protons have a positive electric charge

while the neutrons are electrically neutral. The nucleus makes up almost all of an

atom's mass. Whirling at fantastic speeds around the nucleus are smaller and

lighter particles called electrons which have a negative charge.

3. Using the background information redefine your terms (incorporate charge

when appropriate).

Nucleus:

Protons:

Electrons:

Neutrons:

Figure 1

4. Compare and contrast

your drawing of an atom

to the one on the right

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Figure 2

5. Using the information below (Table 1) complete table 2.

Table 1

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Al Aluminum

27

79

Au Gold

197

29

Cu Copper

64

47

Ag Silver

108

92

U Uranium

238

Table 2

Element Atomic # Atomic

Weight

# of Protons # of

Electrons

# of

Neutrons

Al

Au

Cu

Ag

U

6

C CARBON

12

Atomic Number = represents the number of

Protons as well as the number of Electrons

Chemical Symbol

Chemical Name

Atomic Weight = represents the number of

Protons and the number of Neutrons added

together

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6. Complete the chart using the information provided in Table 3.

Table 3

Element Atomic # Atomic

Weight

# of Protons # of

Electrons

# of

Neutrons

Na 11 23

Cl 17 18

Fr 87 136

Pt 195 78

H 1 1

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2.1 Introduction to the Periodic Table

Background:

The periodic table is arranged in horizontal rows called periods and vertical

columns called groups/ families. The periods have their atomic numbers

increase from left to right, while the groups contain elements having similar

chemical properties. Each group is represented by a Roman numeral and letter.

Letter A represents representative metals, while letter B represents transition

metals. In addition, the figure below highlights several other important factors,

such as ionization energy, metallic properties, and atomic radius.

Organizing the Periodic Table:

The periodic table is arranged into several sections, those being Alkali metals,

Alkali Earth metals, Transition metals, Metals, Metalloids, Non-metals, Halogens,

and Nobel gases.

Directions: Using a copy of the period table and eight different colored pencils

identify each of the sections listed above. Then answer the questions that follow.

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8

9

Characteristics of Periodic Groups

Alkali Metals

Alkali Earth Metals

Transition Metals

Metals

Metalloids

Non Metals

Halogens

Noble Gases

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1. Which pair of elements is in the same period?

2. Which pair of elements is in the same group/family?

3. Which element has the smallest atomic number?

4. If the atomic number of element D is 20, then what is the atomic number of

element R?

(Use your periodic table to answer the remaining questions: 5- 10)

5. This element has the lowest atomic number of any Group 16 element.

6. This element has the most protons of any element in Group 15.

7. This element is in the same family as lead, and it has fewer protons than

sodium.

8. This element has an atomic number that is one greater than platinum.

9. This element has an atomic number that is double that atomic number of

silicon.

10. This element is in Group 1 and has a higher atomic number than chlorine,

but a lower atomic number than bromine.

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Investigating 3.1 ATOMIC PUZZLE Directions: Use the periodic table to determine the element that each lettered question refers

to. Put the element's symbol on the line after each of the clues. When these symbols are

arranged one after the other, they will form a word or phrase. When you discover this word or

phrase, transfer each symbol into the appropriate lettered box. For each problem, a hint is

provided for you.

1. HINT: Produced by the Pancreas?

a. Has 49 positive particles _________

b. Has 6 electrons in its third, and outermost shell ___S______

c. A greenish, radioactive mineral used in the production of

electricity ___U__

d. This element's nucleus contains only 4 neutrons _________

e. The atom with 7 electrons _________

a b c d e

2. HINT: What your teacher will "keep" during a test.

a. Lightest element in the universe _________

b. It has 7 electrons on its 5th, and outermost shell ___i___

c. Has 16 protons _________

d. Has a mass of 127 _________

e. An element used to form water _________

f. Element with 7 neutrons _________

g. Atomic number 39 _________

h. The gas that our blood absorbs during respiration_____

I. Symbol for uranium _________

a b c d e f g h i

3. HINT: They lived in a submarine in 1966

a. Has four protons _________

b. Element with 125 neutrons, in the same family as fluorine _________

c. Roman numeral for "50" ___L______

d. The “Actinoid” element named after "Albert" _________

a b c d

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4. HINT: A typical Irish breakfast

a. Has 104 neutrons _________

b. Symbol for carbon _________

c. Has 7 less electrons than Iron_________

d. Has 39 protons _________

e. Has only a filled inner shell, and half-filled second shell ____

f. An element found in water, and hydrochloric acid _________

g. Has 8 electrons in it's outermost, third shell _________

h. Abbreviation for "Miss" _________

a b c d e f g h

5. HINT: He had an element named after him

a. Symbol of the element used in foil _________

b. The element with 6 neutrons and 5 protons ___B______

c. Contains 67 more positive particles than Hydrogen _________

d. 20th letter of the alphabet _________

e. Letter used to symbolize a negative subatomic particle _E________

f. Has a mass that is 31 greater than krypton _________

g. It has 6 electrons on its 3rd and outermost shell_________

h. Atom whose atomic number is 52 _________

I. Its electron configuration is 2, 8, 18, 18, 3 _________

a b c d e f g h i

6. HINT: Projected onto a screen by a cathode ray tube = TV TUBE

a. Symbol for boron _________

b. The element with two electrons in its 7th shell ___Ra____

c. 10th element of the "lanthanoid" series _DY________

d. It's second, outermost shell has 3 electrons ________

e. Contains a total of 92 protons _________

f. One of the laughing gas elements _________

g. Atomic mass is 12 _________

h. Having no neutrons it is an explosive gas once used in blimps _________

a b c d e f g h

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Investigation 3.2 Laboratory Investigation: Understanding an

Electron’s Behavior

In 1913, the Danish physicist Niels Bohr proposed yet another

modification to the theory of atomic structure based on a curious

phenomenon called line spectra.

Under Bohr's theory, an electron's energy levels (also called

electron shells) can be imagined as concentric circles around the

nucleus. Normally, electrons exist in the ground state, meaning

they occupy the lowest energy level possible (the electron shell closest to the

nucleus). When an electron is excited by adding energy to an

atom (for example, when it is heated), the electron will absorb

energy, "jump" to a higher energy level, and spin in the higher

energy level. After a short time, this electron will spontaneously

"fall" back to a lower energy level, giving off a quantum of light

energy. Key to Bohr's theory was the fact that the electron

could only "jump" and "fall" to precise energy levels, thus emitting a limited

spectrum of light. The animation linked below simulates this process in a

hydrogen atom.

Problem:

How does the emission of spectra of various types of light appear through a

spectroscope?

Materials:

incandescent light

fluorescent light

hydrogen light

helium light

krypton light

spectroscope

high voltage source

Procedure: Part 1

1. Obtain a spectroscope.

2. Look through the spectroscope toward the fluorescent lights in the room.

Observe the color produced.

3. Record the color(s) of the spectra in data table 1 (see example in book).

4. Draw the color(s) of the spectra in the box labeled fluorescent light. Record your

observations by using colored pencils.

5. Repeat steps 2-4 using an incandescent light bulb. Record all observations.

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Part 2

6. Obtain a spectroscope as you instructor inserts a vacuum tube of an element in

the high voltage source.

7. Record the color(s) of the spectra in data table.

8. Draw the color(s) of the spectra in the box labeled fluorescent light. This is the

bright line spectrum. Record your observations by using colored pencils.

9. Repeat steps using the other two vacuum tubes of various elements identified in

the materials.

Observations/Data Table:

Type of Light

Source

Color of gas as

visible without a

spectroscope

Drawing of the continuous and/or bright-

line spectrum

(please used colored pencils for your

drawings)

Fluorescent Light

bulb

Incandescent

light bulb

Helium

Hydrogen

Krypton

Neon

Mercury

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Analysis Questions:

1. White light is made up of all seven parts of visible light. Identify the seven

components.

2. Describe the difference you see between the light bulbs (fluorescent/

incandescent) and the other elements that give off light (vacuum tubes).

3. Of the five different types of light spectra you examined, which ones do you

think are the MOST similar? Explain.

4. Create a prediction regarding how using an emission spectra would be

helpful to scientists?

5. What is the purpose of the high voltage source?

6. Explain the process of how the bright line spectrums are produced in the

samples of light that were examined.

7. How might spectral analysis be useful in astronomy? Think about this question

carefully before you answer. Please provide a detail explanation of this

question.

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Investigation 3.3 Electron Configuration

How does electron configuration relate to the periodic table? The periodic table is organized into Periods (rows), Groups 1-18 (columns) and Blocks (s, p, d and f).

The periodic table below shows the s, p, d and f view:

The arrangement of electrons within the orbitals of an atom is known as the electron configuration. The most stable arrangement is called the ground-state electron configuration. This is the configuration where all of the electrons in an atom reside in the lowest energy orbitals possible. Keeping in mind that each orbital can accommodate a maximum of two electrons, we are able to predict the electron configurations of elements using the periodic table.

Basically, the distributions of orbitals can be laid out in the following fashion (read from the bottom up):

_ 4s _ _ _ 3p _ 3s _ _ _ 2p _ 2s _ 1s

The bottom energy level is level 1 - it has the lowest energy. Each "_" represents an orbital. You can see

that there is 1 orbital for an s subshell. There are 3 orbitals for a p subshell. Each orbital can hold 2

electrons. Therefore, the s subshell can hold 2 electrons and the p can hold 6. As a result, the first

energy level can hold 2 electrons (1s = 2), and the second energy level can hold 8 electrons (2s2p = 2 +

6), etc.

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Parts of an electron configuration:

Energy level - a number (1, 2, 3 and so on)

Sublevel (orbital) - a letter, either s, or p

Number of electrons - a superscript number

Analogy: The energy level is like a driveway with cars in it, the sublevels are the type of cars in parking

lot, and the orbitals are how many seats are in the car.

How to write an electron configuration:

In a neutral atom, the number of electrons equals the number of protons of the atom. When the electrons

fill the orbitals, they occupy the lowest energy orbitals that are available.

For example, hydrogen is atomic number 1 (has 1 proton). The one electron that it has occupies the

lowest orbital, which is 1s. To write its electron configuration, it would be 1s1. In an orbital diagram, it

would simply be a line with one up arrow in it, which represents the 1s orbital:

H: 1s1

Practice: Write the electron configuration for the following elements-

1s2 2s2 2p6 3s2 3p6 4s2

He Li

Cl P Ca C O F B Ne

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Investigation 3.4 Describing Valence Electrons using Lewis Dot

Structures

Background:

American chemist Gilbert Lewis developed another method for depicting an

atom’s electron configuration. Lewis focused on the valence electrons in his

representations of elements. Valence electrons are all of the electrons in the

outer orbital/shell of an element.

1. Based upon the above information how many valence electrons do the

following elements have:

a. B

b. Mg

In a Lewis dot structure the valence electrons are represented as dots placed

around the element symbol.

2. Draw what you think the Lewis dot structure looks like for:

a. B

b. Mg

Check your predictions by sharing with the class.

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Draw the Lewis-Dot Structure for the following atoms. Review the following

bullets before you begin.

Around the element’s symbol use dots to represent each valance

electron

The dots should be spread over the four sides. Dots are not paired until all

sides have at least one dot.

Place the initial dot above each symbol and then proceed to arrange the

rest in a clockwise direction.

The number of valance electrons is equal to the group number (Roman

numeral).

1. B

2. Ne

3. Li

4. He

5. C

6. P

7. S

8. Mg

9. H

10. F

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Investigation 3.5: Understanding the formation of bonds and ions. The octet rule says that atoms tend to gain, lose or share electrons so as to have eight

electrons in their outer electron shell.

When atoms form ions they seek to obtain a stable electron

configuration (8 valance electrons—Lewis-dot structures). As a result,

they will either attempt to gain electrons and or lose electrons to

become stable. In either situation they will do whatever is the easiest. The

protons (+) in the nucleus of an atom remain unchanged by ordinary chemical

reactions, but atoms readily gain or lose electrons (-). When electrons are removed

from or added to a neutral (non-charged) atom, a charged particle called an ion is

formed. If the atom gains electron(s), its net charge becomes negative. A negative

ion is called an anion. If the atom loses electron(s), its net charge becomes positive. A

positive ion is called a cation. Gaining or losing electrons allows every energy level in an

ion to hold only a "stable" number of electrons, namely, 2, 8, 18 or 32.

How many electrons are in the following ion

a. Al 3+ _________

b. Te 2- _________

c. Si 4+ _________

d. Sb 3- __________

e. Cs 1+ _________

f. He _________

g. C4- _________

h. Sr 2+ _________

i. F 1- _________

j. Bi 3- ________

II. According to the Octet Rule, what do these ions need to do in order to become

stable? Indicate how many electrons they would gain or lose.

a. Calcium: Ca ____

b. Boron: B _________

c. Chlorine: Cl _________

d. Indium: In __________

e. Sulfur: S _________

f. Phosphorus: P _________

g. Sodium: Na _________

h. Tin: Sn _________

i. Argon: Ar _________

j. Hydrogen: H _________

III. Are the following atoms anions (a negative ion), cations, both, or neither.

a. Magnesium:___________ d. Neon: ______________

b. Gallium: _____________ e. Carbon:_____________

c. Rubidium:____________ f. Iodine: ______________

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Investigation 4.1: Replacement Reaction

Procedure:

Safety- During this lab you will be wearing googles, gloves and should make

sure that you do not spill any solutions on yourself as there is a strong

possibility of staining clothes.

1. Obtain a piece of metal from your teacher and mass it. Record this

mass in your notes.

2. Obtain a beaker for the mystery solution.

3. Obtain 0.9 grams of the mystery powder and 50 milliliters of water from

the tap and dissolve the mystery powder in the water using a stirring

rod. To avoid stains, be sure to rinse the stirring rod with distilled water

before setting it down.

4. Place the piece of metal in the mystery solution, making sure not to

splash any of the solution.

5. Make note of any reaction that takes place in the beaker.

6. After 10 minutes, gently shake the metal to remove the accumulated

solid. Then place the metal on a paper towel on your lab bench to

allow it to dry.

7. Once dry, mass the metal again and record this value in your notes.

Filtering Procedure:

1. Obtain a piece of filter paper and initial the paper with a pencil. Then,

mass the filter paper, recording this in your data table.

2. Fold the filter paper using the following visuals as a guide:

3. Place the filter paper in a funnel and place another beaker below the

funnel.

4. Pour the solution from your beaker through the filter paper, allowing all

of the solution to transfer to the bottom beaker.

5. Take your filter paper and let it dry. Once dry, mass the filter paper

and precipitate. Subtract the mass of the filter paper to get a final

value for amount of precipitate. Record all data in your data table.

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Analysis Questions:

1) Describe what happened in the beaker when you placed the metal in

the mystery solution?

2) Provide quantitative evidence to support what you described in the

answer above.

3) Hypothesize what you think occurred when the metal and mystery

solution were combined.

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Investigation 5.1 Ionic Bonds

In ionic bonding, electrons are completely transferred from one atom to

another. In the process of either losing or gaining negatively charged

electrons, the reacting atoms form ions. The oppositely charged ions are

attracted to each other by electrostatic forces, which are the basis of the

ionic bond.

The following bullets characterize ionic bonding:

Transfer of electrons

Form between metals and nonmetals

Are crystalline solids

Dissolve easily in water

Conduct electricity in solution

Complete the chart for each element.

Element # of Protons # of Electrons # of Valance Electrons

Sodium

Chlorine

Beryllium

Fluorine

Lithium

Oxygen

Potassium

Directions:

Write the symbol for each element

Using a different colored pen/pencil create a Lewis dot structure for each

element

Draw an arrow(s) indicating the transfer of the electron(s)

Determine the charge of each ion and write the formula

Make sure the sum of the bond is zero

1. Potassium + Fluorine

2. Magnesium + Iodine

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3. Sodium + Oxygen

4. Sodium + Chlorine

5. Calcium + Chlorine

6. Aluminum + Chlorine

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Investigation 6.1Covalent Bonds

Covalent Bonding

Covalent bonding occurs when atoms share electrons. As opposed to ionic

bonding in which a complete transfer of electrons occurs, covalent

bonding occurs when two (or more) elements share electrons. Covalent

bonding occurs because the atoms in the compound have a similar

tendency for electrons (generally to gain electrons).

The following bullets characterize covalent bonding:

Share electrons

Generally occur between two nonmetals

Complete the chart for each element.

Element # of Protons # of Electrons # of Valance

Electrons

# of Electrons to

Fill Outer Shell

Hydrogen

Oxygen

Chlorine

Carbon

Fluorine

Helium

Lithium

Directions:

Write the symbol for each element

Using a different colored pen/pencil create a Lewis dot structure for each

element

Rearrange the electrons to pair up electrons from each atom.

Draw circles to show the sharing of electrons

Write the chemical formula for each molecule

1. Hydrogen + Hydrogen (Diatomic Element)

2. Hydrogen + Oxygen

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3. Chlorine + Chlorine (Diatomic Element)

4. Oxygen + Oxygen (Diatomic Element)

5. Carbon + Oxygen

6. Carbon + Hydrogen

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Investigation 7.1: Lab: Ionic vs. Covalent Bonds

Hypothesis: (Based on the properties of ionic and covalent compounds make a

hypothesis as to which type of bond each compound will have)

Compound Hypothesis Compound Hypothesis

Gelatin Magnesium Oxide

Sugar Detergent

Salt Cornstarch

Baking Soda Copper Chloride

Procedure:

1. You will need safety glasses, lab apron, and gloves.

2. Label 8 pieces of paper towel (good size) with the names of the compounds.

3. Place 5 grams of the compounds on each piece of paper (do ONE at a time).

4. Observe and WRITE the description of the compounds on the data chart.

5. Place 200 ml. of water in a beaker.

6. Add about ½ of the compound to the water and stir with a stirring rod for 10

seconds and observe. If the compound dissolves write SOLUBLE in the data chart

under “Solubility”. If it does not dissolve, write INSOLUBLE.

7. Place the leaders of the conductivity tester (as shown below) into the mixture.

Observe whether or not the bulb lights. If it does, write YES under

“Conductivity” in the data chart. If it does not light up write NO.

8. Rinse and dry the beaker.

9. Repeat the steps for each of the compounds.

10. Be sure to follow the teacher’s directions for how to dispose of the mixtures. DO

NOT just dump them into the sink.

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Data:

Compound Description Solubility Conductivity Bond

Gelatin

Sugar

Salt

Baking Soda

Magnesium Oxide

Detergent

Cornstarch

Copper Chloride

Analysis Questions:

1. Some characteristics of many ionic compounds include solubility in water and

the ability to conduct electricity. On the basis of these two properties, which

compounds appear to have ionic bonds (please indicate in last column “Type

of Bond” in data chart)?

2. Water solutions of covalent compounds do not conduct electricity. Based on

this property, which compounds that you tested would you classify as

covalent compounds (please indicate in last column “Type of Bond” in data

chart)?

3. Explain how ionic and covalent compounds are different. USE

EVIDENCE FROM YOUR LAB.

4. Did all of the compounds that conducted electricity show the same amount of

conductivity? How can you tell (from the lab…)?

5. Solid table salt does not conduct electricity. Why do you think dissolving

table salt in water allows the salt to conduct electricity?

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Unit 2 Study Guide-Atoms, Bonds, ions, and their

properties

Key Terms:

You must be able to define and explain each of the following terms below

Atom

Molecule

Element

Proton

Neutron

Electron

Nucleus

Ionic bonding

Covalent bonding

Isotope

Ion

Cation

Anion

Lewis Dot

Atomic number Atomic weight group/family

Period Octet Rule Valence Electron

Application:

You must be able to evaluate and utilize the periodic table in order to:

Decipher the number of proton, electrons, neutrons, atomic

number, and atomic weight.

Evaluate whether the element is a Alkali Metal, Alkali Earth Metal,

Transition Metal, Metal, Metalloide, Non-Metal, Halogen, or a Noble

gas

In terms of characteristics, be able to identify and explain the

similarities and difference between and among the eight different

groups

Situate the electrons into their respective orbital using the electron

configuration model and the Lewis-Dot model.

Interpret and explain ions according to the Octet Rule, example

Ca 20, protons 20 electrons, and 20 neutrons. According to the

Octet Rule it is simpler for Ca to give away 2 electrons than gain

eight. Therefore, Ca becomes Ca2+ with electron configuration of

2-8-8.

Determine whether an ion is a cation, anion, or neither.

Draw and depict a covalent and ionic bond

Graphing/ Short Response:

Be able to organize and synthesize data

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Unit Three Key Terms

Directions: Define the following terms using either your book or

the internet. Please attempt to make the definition clear,

concise, and simple.

1. Atom

2. Molecule

3. Element

4. Proton

5. Neutron

6. Electron

7. Nucleus

8. Ionic bonding

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9. Covalent bonding

10. Ion

11. Cation

12. Anion

32

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