In this lesson students learn about the trends of the Periodic Table by doing an inquiry activity, taking notes, watching videos, and doing practice questions. Within this unit students already learned about electronegativity in terms of bonding types in lesson 6: Metallic and Covalent Bonds.
To begin the lesson I have students begin to figure out the trends for ionization energy, electronegativity, and atomic radius through an inquiry activity. In this section of the lesson I go through the trends with students. My goal is for students to learn what the trends are but also understand why the trends occur.
The lesson begins on slide 6 of the PowerPoint and students fill in their notes graphic organizer while I present the information. I begin the lesson by going over effective nuclear charge slides as this is an important aspect of the trends. To use an analogy I talk about Tommy the Trojan our school mascot and how his charm is the effective nuclear charge with those students closest to him being most effected by his charm. I also use the analogy of students in a classroom and those closest to the teacher are watched more closely.
I then go into the three major trends for atomic radius, ionization energy, and electronegativity with students filling out information in a table on their graphic organizer. This is a copy of the filled in notes that are expected from students. To reinforce the trends I have students perform a practice worksheet. The worksheet contains two sections. This year I decided to have students do the first section of the worksheet with partners. I did this because I could tell that after doing the trends activity and notes that students needed to get up and move around a bit.
Because I did not initially make the worksheet as a partner practice paper I had students write at the top of their paper and find partners.
I then assigned them questions to work on with their partners. For additional information on how I use partner appointments in my classroom please look at the Partner Appointments reflection in Unit 1 Lesson 7: Dimensional Analysis. The questions for trends are tricky and a lot of students want to just right the trend, but not explain why the trend occurs.
To help students I continuously remind them to look back at their notes for help with the rationale for the trend.Tawan arb dao ep 1 eng sub
This is a copy of the answer key for the first section with how I added in partner appointments at the top and the questions I had students perform with their partners.Any chemistry instructor who might have these two files should contact. The computer simulation was developed by Prof. A student activity tutorial accompanies this computer simulation. The activity was developed by Prof. By continuing to view the descriptions of the demonstrations you have agreed to the following disclaimer.
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The custom demos section of the website is used by UO chemistry instructors to schedule demonstrations that are not listed in the database. Do not proceed to schedule a custom demo unless you have already conferred with the lecture demonstrator about it. Skip to main content. UO Home Dept Index. Periodic Trends Computer Simulation. Curriculum Notes.
Schedule This Demo. Attached Files Periodic Trends Activity. Please read the following disclaimer carefully By continuing to view the descriptions of the demonstrations you have agreed to the following disclaimer. Custom Demos Warning The custom demos section of the website is used by UO chemistry instructors to schedule demonstrations that are not listed in the database.The periodic table of elements has a total of entries.
Elements are arranged in a series of rows periods in order of atomic number so that those with similar properties appear in vertical columns. Elements in the same period have the same number of electron shells; moving across a period so progressing from group to groupelements gain electrons and protons and become less metallic.
This arrangement reflects the periodic recurrence of similar properties as the atomic number increases. For example, the alkali metals lie in one group Group 1 and share similar properties, such as high reactivity and the tendency to lose one electron to arrive at a noble-gas electron configuration.
Modern quantum mechanics explains these periodic trends in properties in terms of electron shells. The filling of each shell corresponds to a row in the table. In the s-block and p-block of the periodic table, elements within the same period generally do not exhibit trends and similarities in properties vertical trends down groups are more significant.
However, in the d-block, trends across periods become significant, and the f-block elements show a high degree of similarity across periods particularly the lanthanides. If we examine the physical state of each element, we notice that on the left side of the table, elements such as lithium and beryllium are metallic solids, whereas on the right, nitrogen, oxygen, fluorine, and neon are all gases.
This is because lithium and beryllium form metallic solids, whereas the elements to the right form covalent compounds with little intermolecular force holding them together. Therefore we can say that, in general, elements tend to go from solids to liquids to gases as we move across a given period. However, this is not a strict trend. As you move across a period in the periodic table, the types of commonly encountered bonding interactions change. For example, at the beginning of Period 2, elements such as lithium and beryllium form only ionic bonds, in general.
Moving across the period, elements such as boron, carbon, nitrogen and oxygen tend to form covalent bonds. Fluorine can form ionic bonds with some elements, such as carbon and boron, and neon does not tend to form any bonds at all. Another physical property that varies across a period is the melting point of the corresponding halide. A halide is a binary compound, of which one part is a halogen atom and the other part is an element or radical that is less electronegative or more electropositive than the halogen, to make a fluoride, chloride, bromide, iodide, or astatide compound.
Many salts are halides; the hal- syllable in halide and halite reflects this correlation. All Group 1 metals form halides that are white solids at room temperature. The melting point is correlated to the strength of intermolecular bonds within the element.
First, we must analyze compounds formed from elements from Groups 1 and 2 e. To develop an understanding of bonding in these compounds, we focus on the halides of these elements.
The physical properties of the chlorides of elements in Groups 1 and 2 are very different compared to the chlorides of the elements in Groups 4, 5, and 6. All of the alkali halides and alkaline earth halides are solids at room temperature and have melting points in the hundreds of degrees centigrade. The non-metal halide liquids are also electrical insulators and do not conduct electrical current. In contrast, when an alkali halide or alkaline earth halide melts, the resulting liquid is an excellent electrical conductor.
This tells us that these molten compounds consist of ions, whereas the non-metal halides do not. This again demonstrates the type of bonding that these compounds exhibit: the left-most elements form more ionic bonds, and the further-right elements tend to form more covalent bonds.
The physical properties notably, melting and boiling points of the elements in a given group vary as you move down the table. In chemistry, a group is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table, including the d-block elements but excluding the f-block elements.
A physical property of a pure substance can be defined as anything that can be observed without the identity of the substance changing. The observations usually consist of some type of numerical measurement, although sometimes there is a more qualitative non-numerical description of the property.Now that we have an understanding of the electron configurations of the periodic table we are ready to tackle periodic trends in atomic properties.
This is a very important section as it will allow us to predict behavior of atoms of the various elements. For example, if two ions have the same charge, the smaller ion forms a stronger bond because it can get closer to the counter ion.
So knowing the radius of ions can help us predict ionic bond strengths. In simpler terms, it can be defined as something similar to the radius of a circle, where the center of the circle is the nucleus and the outer edge of the circle is the outermost orbital of electron.
As you begin to move across or down the periodic table, trends emerge that help explain how atomic radii change. With the exception of the 1s orbitals the nuclear charge the valence shell electrons feel is not the full nuclear charge Zbut a reduced charge due to shielding caused by the negative charge of the inner shell core electrons. Now determining the size of an atom is actually impossible because according to the Heisenberg Uncertainty Principle we never really know where an electron is, and is going at the same time.
But if we look at two bonded atoms we can approximate the size by looking at the distances between the atoms, that is, by looking at bond lengths we can get an approximation of the size of the atoms.Statistica inferenziale 1
We will take two approaches based on the type of bond. When a covalent bond is present between two identical atoms, the covalent radius can be determined. When two atoms of the same element are covalently bonded, the radius of each atom will be half the distance between the two nuclei because they equally attract the electrons. The distance between two nuclei will give the diameter of an atom, and the covalent radius is half this diameter.
Some compounds like the noble gasses do not form bonds and you can cool them down until they form a solid, and measure the difference between them, which is know as the Van der Waals atomic radius.D battery in vinegar in a glass
These are not necessarily equal. Ionic radius are more complicated because the anion and cation are often of vastly different sizes. Now as we shall see, anions are larger than neutral molecules because they gain electrons, and metals are smaller than neutral atoms because they lose electrons. So it is very common for the anion to be larger than the cation. As we shall see in Chapter 12 section What is important to note here is that you can not use the internuclear distance as you can with covalent bonds, because the ion sizes are not equal.
At first glance you would think you could predict metallic radii the same way you do for covalent, that it would be one half the distance between two metal atoms.
As we will learn in Chapter 12 section There are two general periodic trends in atomic radii. The radii increase as you go down a group, and decrease as you go across a period.To get the best possible experience using our website, we recommend that you upgrade to latest version of this browser or install another web browser. Network with colleagues and access the latest research in your field.
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Periodic Table Trends
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7.3: Atomic Properties and Periodic Trends
If you do not respond, everything you entered on this page will be lost and you will have to login again. Don't show this again! Careers Launch and grow your career with career services and resources.Need Help? Last updated March 25, In this simulation, students will investigate several periodic trends, including atomic radius, ionization energy and ionic radius.
Through the use of this simulation students will have the opportunity to examine atomic data as well as visually compare and interact with select elements from the periodic table.
In this investigation you will examine several periodic trends, including atomic radius, ionization energy and ionic radius. You will be asked to interact with select atoms as you investigate these concepts.
Using your computer, tablet or mobile device, navigate to the website: teachchemistry. You should see the picture below on your screen. Write the symbols and atomic number for each of the elements that you chose below:.
Do these values support your answer for part b? Reset the selected data using the reset symbol. Again choose a second element to compare from the same group. Do these values support your answer for part e? Give suggestions for why you think this trend exists based on your interaction with the elements. Again choose a second element to compare from the same period.
List them below:. Using the simulation, check your predicted answers to see if you are correct! Describe what happened.
Were you successful? Did the atom seem to have a strong hold on the electron? Attempt to ionize this atom by pulling a valence electron from the electron shell. What trend in ionization energy do you observe for elements in the same period based on the data in the graph? For example, do atoms with larger atomic numbers have greater ionization energy than atoms with small atomic numbers in the same group?
Navigate back to the main page, and reset the data using the reset symbol. How do their ionization energy values compare? Does this data support your prediction from part h? Explain your choice referencing both the atomic model and subatomic particles:. What happened to the electron shell where this valence electron was located? How is the change in subatomic particles related to the size of the ion?
What is the value? Is this value larger or smaller than the value for the atomic radius in part a? Justify your prediction with scientific reasoning.Choose your answers to the questions and click 'Next' to see the next set of questions. You can skip questions if you would like and come back to them later with the yellow "Go To First Skipped Question" button.Upload video to identify song
When you have completed the practice exam, a green submit button will appear. Click it to see your results. Good luck! Because the outermost electrons get closer and closer to the atom's nucleus, the attraction between the nucleus and bonding electrons decreases.
Because the outermost electrons get closer and closer to the atom's nucleus, the attraction between the nucleus and bonding electrons increases. Because the outermost electrons get farther and farther away from the atom's nucleus, the attraction between the nucleus and bonding electrons increases. Because the outermost electrons get farther and farther away from the atom's nucleus, the attraction between the nucleus and bonding electrons decreases.
This group contains poisonous, very reactive elements often used in bacteria killing or household cleaners. Premium members get access to this practice exam along with our entire library of lessons taught by subject matter experts.
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Periodic Trends Guided-Inquiry Activity
Exam Instructions: Choose your answers to the questions and click 'Next' to see the next set of questions. Answered 0 of 30 questions. Start Exam. Page 1 Use this material to answer questions 1 through 5. Question 1 1. How many valence electrons does silicon have? Use this material to answer questions 1 through 5. Question 2 2. How many valence electrons does krypton have?
Question 3 3. Which of the following elements has the same number of valence electrons as aluminum? Question 4 4. In which energy level is the outer electron in rubidium located?
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