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Periodicity

Periodicity refers to the recurring trends seen among the elements of the periodic table. By arranging the elements in order of increasing atomic number trends can be easily interpreted. It was actually because of periodicity which made the completion of the modern periodic table possible – as it could help predict the properties and atomic mass of the missing elements.

The main periodic trends are ionization energies, atomic radius, electronegativity and electron affinity. Ionization energy refers to the amount of energy required to move one valence electron of an element in the gaseous state. Although atomic radius can’t be measured directly, it can be inferred by halving the distance between the centers of two atoms engaged in a covalent bond. Electronegativity is the measure of an atom's ability to form a bond; while electron affinity is a measure of an atoms attraction towards a particular electron.

With the arrangement of the periodic table in rows and columns based on the number of main electron shells and number of valence electrons, respectively, these trends can be observed based on these electron characteristics. If an electron is further away from the nucleus then then the ionization energy is small as it is easier to remove; the atomic radius is large as the electron is far from the nucleus’ the electronegativity would depend on the number of protons; while the electron affinity would also be low as the attraction is not very strong. Thus, understanding periodicity is extremely important for predicting the properties of elements.

Group Characteristics in the Periodic Table

Research Question: You are a group of young tutors teaching a group of 12 graders preparing for university entrance; your goal is to teach them periodic table group characteristics with emphasis on specific elements of each group. You will include in your body of the report: Hydrogen, Alkali metals, earth alkali metals, ra

Ionization Energies

Why does lithium have a larger first ionization energy than sodium? b. the difference between the third and fourth ionization energies of scandium is much larger than the difference between the third and fourth ionization energies of titanium. Why? Please explain in detail.

First and Second Ionization Energy for Mg and Na

The first ionization energy for Mg (738 kJ/mol) is higher than that for Na (469 kJ/mol) but the second ionization energy for Na (4562 kJ/mol) is much larger than that for Mg (1451 kJ/mol). Explain this observation.

Graphing Atomic Radii and Ionisation Energy for Elements

Graphing of atomic radii and ionisation energy for elements 3 - 30 in the periodic table/Questions are answered on the trend in atomic radii for alkali metals and alkali earth metals down a period as atomic number increases and how the ionisation energy trends across a period. See the attached file.

9 questions about energies and periodic trends

1. Which process is energetically favorable: a) adding an electron to K to form K-; or b) losing an electron to form K+? Explain. 2. Using Slater's rules, determine the effective nuclear charge for a 4s and 3d electrons in vanadium and V2+. Discuss the relative sizes of the atomic/ionic radii based on your result. 3.

Group Trends: Metallic Character and Increasing Acidity

1) For each of the following pairs, which element will have the greater metallic character. a) Li or Be b) Li or Na c) Sn or P d) Al or B 2) Arrange the following oxides in order of increasing acidity. Rank from least acidic to most acidic. To rank items as equivalent, overlap them. CO2, P2O5, SiO2, CaO, SO3, Al2O3

Ionization Energy and Electron Affinity

1) Which ion is smaller? F- or Na+ 2) For isoelectronic ions, how are effective nuclear charge and ionic radius related? As nuclear charge (Z) increases, ionic radius increases or As nuclear charge (Z) increases, ionic radius decreases. 3) For the elements with the electron affinities given in the table in the introduc

Ionization Energies and Electron Affinities: Example Problems

1) Based on position in the periodic table and electron configuration, arrange these elements in order of decreasing I1. Rank the elements from highest to lowest ionization energy. To rank items as equivalent, overlap them. K, O, Na, P, Al, S 2) Examine the following set of ionization energy values for a certain element. How

Ionic Radii: Example Questions

1) Consider the relation shown for the radius of the species indicated. Which choice presents a valid comparison? a) Na>Na+ b) Ca2+>Ca+ c) Cl>Cl- d) S->S2- 2) The following ions contain the same number of electrons. Rank them in order of decreasing ionic radii. Rank from largest to smallest radius. To rank items as equiv

Effective Nuclear Charge and Sizes of Atom and Ions: Example

1) Using only the periodic table, arrange each set of atoms in order of increasing radius. Rank from smallest to largest. To rank items as equivalent, overlap them. a) Be, Mg, Ca b) Ge, Br, Ga c) Si, TI, Al 2) Rank the following elements in order of decreasing atomic radius.Rank from largest to smallest radius. To ra

Periodicity and Chemical Bonding

1. How does ionization energy change as one proceeds a) across the periodic table in a horizontal row b) down the periodic table in a vertical group Give an explanation for each answer. 2. Arrange the following atoms or ions in order of increasing radius: Cl, S2-, K, K+, O Give an explanation for the position o

ELECTRON CONFIGURATIONS and PERIODIC TRENDS

1. Briefly explain what is wrong with the following orbital diagrams. 2. Briefly explain what is wrong with the following electron configurations. 3. (Periodic Trends: Atomic Radius) a. As you go down a column in the periodic table, atomic size increases. Explain why in terms of principle quantum number and effect

Ionization Energy and Electron Affinity Problems

Of the choices below, which gives the order for first ionization energies? a. Ar > Cl > S > Si > Al b. Cl > S > Al > Si > Ar c. Al > Si > S > Cl > Ar d. Cl > S > Al > Ar > Si Which of the following has the largest second ionization energy? a. Na

Periodic Table Trends

Electron affinity (E.A.) generally becomes more negative across a period, from left to right. However, the E.A. of phosphorus does not follow this trend and is less negative than that of silicon. Using orbital diagrams showing the valence electronic configuration of Si and P, explain why this exception occurs. Clearly justify yo