BASIC CHEMISTRY 1 MID TEST SEMESTER

Mid Semester Test

Course             : Basic Chemistry
Credit            : 3 SKS
Lecturer           : Dr. Syamsurizal

( EKA MUHARYANI SIREGAR – RSA1C112008 )

1.     Pure substance X solid at room temperature. If the subtance is heated to 230⁰Cis melted gradually.if then cooled to room temperature,the liquid can not befrozen.

a.      Is it possible X of an element or a compound. Explain it!

b.      Does it a chemical change occured? If so can it be said to undergo an endotherm changing, based on the information provided?

c.       Can it be said that the liquid  is an element,based on the  information provided.

2.       When a candle that weight 10 g is burned  in oxygen, carbon dioxide and water vapor formed by combustion the weight  more than 10g. Was this case match with the law of conversation of mass. Explain it!

3.       When carbon burns in oxygen under limitid number,it will form two gaseous compounds. Suggest the way to differentiate the two compounds with one another.

4.       After mendeev compiled the periodic tabl, he concluded that the atomic weights of certain elements was wrong ruling, and this conclusion was apparently correct. How medeleev was able  to predict that several atomic weights were wrong? Why his predictions are  not always right. Explain!

5.       When an aqueous solution of merkuri chloride is added to an aquades solution  of silver nitrate, a white solid froms. Identify the white solid and write the balanced equation for the reaction that occurs

ANSWER
1. a. I think X is a compound, because the compound is a single substance that consists of several elements are intertwined. The compound is formed from at least two different elements. Although made up of different elements, but the compound is still considered a single substance, because the properties of the elements that make it up can not be found in the compound. The compounds can also be broken down into its constituent elements by a chemical reaction. All compounds will break down into smaller compounds or individual atoms when heated to a certain temperature (called the decomposition temperature). Compounds have properties that are different from their constituent elements. DAPT compound just broken into its constituent elements by chemical reactions. In the same conditions, the compound may have a different form of the constituent elements. For example, the reaction between hydrogen gas and oxygen gas to form water compounds cair.Ciri typical compound is that it has a fixed constituent mass ratio.

b. yes, because the chemical reactions to form new compounds change, the sustainability of a chemical reaction can be observed from the changes, those changes include color changes, changes in temperature, precipitation formation, and the formation of gas, and the compound X was found that the temperature change kimia.ada reactions can be said that changes in said transformed into chemical reactions endothermic reaction is a chemical reaction that absorbs or receives heat, heat transfer occurs from the environment to the system so that the ambient temperature dropped and became cooler and because the fluid has undergone endothermic reaction so room temperature does not affect the shape of the liquid due to the possibility of liquid can freeze at a certain temperature.

c. yes, because the elements are the basic parts of the underlying things (blocks of). namely simple substances (singular) who is considered as the composition of the universe (such as air, soil, water, fire) section (which is important, required) dr whole even bigger and substances that can not be broken down into simpler substances that (which is not complex). 

2. Yes, I think that according to the law of conservation of mass as the Law of Conservation of Mass formulate Lavoisier (Lavoisier) which reads:In a chemical reaction, the mass of substances before and after the reaction is the same. While the candles burning with oxygen which weigh 10 grams are changed during the combustion becomes heavier 10gr, while the mass was not changed despite burnt / in the heat but the change from burning candles are not mass weight.

3. if the carbon in the fuel with oxygen gas will produce compounds, namelyC (S) + O2 (g) = CO2C (S) + 1/2O2 = COhow to distinguish them is by doing experiments that the combustion in the form of a compound which is then get the difference as follows:oxygen is relatively stable because oxygen can react with other elements either metallic or non-metallic carbon whereas usually only react with oxygen and hydrogen, and oxygen is used for breathing while the carbon compounds that are toxic if inhaled.Carbon monoxide gas can be harmful if inhaled excess carbon dioxide but may be required by the plants.

4. In his research, he made some sequence of chemical elements known to establish a set of cards, where each card contains every element and chemical properties. He discovered the periodic law when drawing cards, which are arranged in a sequence according to increasing atomic mass, and its properties are repeated periodically. There's even a little excess on the periodic table are arranged, namely the recognition of new elements as yet unknown to the masses he discovered the periodic table such as gallium, germanium, and scandium. He provides space for the elements and even estimate how the mass of the atom, and how their chemistry nanti.Mempertimbangkan Mendeleev periodic law can not tolerate any exceptions for any deviation from the order of the atomic weight of the element. For example, in the case of the elements tellurium and iodine, he predicted that the atomic weights of these elements are not appropriate because the values ​​provided indicate the order opposite to what is determined by the chemical nature. More specifically, tellurium showed a higher atomic weight in accordance with the measured values, and then the order of elements based on this feature will put the tellurium in the same chemical group as fluorium, klorium, and bromium, where it is not included in the periodicity of the used chemicals. Apparently, according to Mendeleev reversal sequence is correct, but the reason that is used inappropriately. It really weighs close to atomic accuracy, but the order of elements is better placed based on the number of atoms of each element of the sequence unsur.Skema enhanced by the work of Moseley in 1912, and the accuracy of the main complication lies in the mixture of isotopes occurring in the chemical elements the most important.Mendeleev periodic believe that the law could be used to predict the existence of new elements some properties of the compound, as well as to correct the atomic weights of some elements already known. However, historians and writers too stressed the prediction of chemical in its legal aspects. Mendeleev ability to accommodate elements that are known to have a positive contribution in the periodic system.

5.   HgCl2 + 2AgNO3=Hg (NO3) 2 + 2AgCl

AgCl precipitate formed in the gray-white colored solid nitrate

THE STRUCTURE OF ATOM AND PERIODIC TRENDS

Periodic Trends

Introduction

Electrons are held in an atom or ion by the electrostatic attraction between the positively charged nucleus and the negatively charged electrons. In multi-electron species, the electrons do not experience the full positive charge of the nucleus due to shielding by electrons which lie between the electron of interest and the nucleus. The amount of positive charge that actually acts on an electron is called the effective nuclear charge.

Effective Nuclear Charge

The concept of effective nuclear charge (Z*) is important to understanding periodic properties. The effective nuclear charge is that portion of the total nuclear charge that a given electron in an atom experiences. This is equal to the atomic number (Z) minus the amount (σ) that other electrons in the atom shield the given atom from the nucleus.

Z* = Z-σ

Example: Lithium has three protons and an electron configuration of 1s²2s¹. The electron in the 2s orbital is shielded from the full attraction of the protons by the electrons of the 1s orbital (Figure 1). Thus, Z* felt by the 2s electron should be one rather than three. However, lithium's 2s electron does not behave as if it is experiencing exactly a +1 charge (Z* is actually about 1.3 charge units). This can be explained by the fact that the 2s orbital has two maxima in its radial probability function (Figure 1), and the lesser maxima penetrates within the maximum of the inner 1s electron. Although lithium's 2s electron spends most of its time in the outer lobe of that orbital feeling a nuclear charge of +1, some of the time it is inside the 1s orbital experiencing the full nuclear charge of +3. Thus, Z* is somewhat greater than +1.

Slater's Rules for Determining σ

In 1930, J.S. Slater formulated the following set of empirical rules for determining the values of the shielding constant σ.

Slater's Rules

Write out the electronic configuration of the element and group the orbitals in the following order: (1s)(2s, 2p)(3s, 3p)(3d)(4s, 4p)(4d)(4f)(5s, 5p)........ To establish the screening constant for any electron, sum up the following contributions: Electrons in groups outside (to the right) of the one being considered do not contribute to the shielding. Electrons in the same group contribute 0.35 to the shielding (except the 1s group, where a contribution of 0.30 is used For s or p electrons being observed, each electron in the (n-1) shell contributes 0.85 to the shielding and each electron in the (n-2), (n-3), ... shells contribute 1.00 to the shielding For d or f electrons being observed, each electron in an underlying group contributes 1.00 to the shielding.

Example

Example: Calculate Z* for a 4s and a 3d electron in Zn Determine the electron configuration for Zn (1s)2(2s, 2p)8(3s, 3p)8(3d)10(4s)2 For a 4s electron: Establish the screening constant for the 4s electron σ = (1 x 0.35) + (18 x 0.85) + (10x1.00) = 25.65 Calculate the effective nuclear charge Z*= Z-σ = 30-25.65 = 4.35 For a 3d electron: Establish the screening constant for the 3d electron Calculate the effective nuclear charge From this example, you can see that the 3d electrons experience a much greater positive charge than the 4s electron and would be held more tightly. Thus, the 4s electrons will be the first removed when Zn is ionized.

PROBLEMS

Using Slater's rules, calculate a value for the effective nuclear charge felt by (a) an electron being added to the 3s orbital of a neon atom and (b) an electron being ionized from the 2p orbital of the neon atom. Comment on your results relative to the stability of the electron configuration of the neon atom. Calculate Z* for the valence electrons in the atoms Li to Ne using (a) the assumption that σ equals the number of inner-shell electrons and (b) Slater's rules. Plot both sets of results on the same graph and discuss. Recall why the energy of an ns orbital is less than that of an np orbital. Use this information to discuss the assumption that these orbitals are always considered as a group (ns, np) in Slater's rules. Plot of the probability of finding 3s, 3p, 3d and 4s electrons as a function of the radial distance from the nucleus can be viewed here. Discuss these probabilities relative to rules 2c and 2d of Slater's rules. You can access a spreadsheet for calculating effective nuclear charges here. If you need assistance in using Excel for plotting data, try this tutorial.

The Periodic Table

You have used the periodic table throughout your study of chemistry. Read more about the periodic table here. Mendeleev was one of the early chemists to recognize that the properties of the elements were periodic in nature. Read from Mendeleev's original publication. To get a flavor of what it would be like to derive the periodic table, try this simulation. It is like discovering the pieces of a jigsaw puzzle and putting them together.

Periodic Trends

In the remainder of this module, you will be analyzing the periodic trends that exist among the elements. Start your investigation by viewing this movie on periodic trends. To view this movie, you may need to download and install the Real Player.

Atomic Radius

There are several ways to define the atomic radius of an atom:

• covalent radius (r_cov), the half-distance between the nuclei of two atoms joined in a covalent bond
• van der Waals radius (r_vdw), the half-distance between the nuclei of two atoms of neighboring molecules
• metallic radius (r_met), for metallic elements, the half-distance between the nuclei of two neighboring atoms in the solid metal

The difference between covalent radius and van der Waals radius is shown in Figure 2.

PROBLEMS

Plot the values of radii vs. atomic number for the Group 1A elements and the Period 2 elements. you will find the data that you need in this Excel spreadsheet of physical property data for the elements (right clicking on the link will allow you to save the file). Use the concept of effective nuclear charge to rationalize the trend in radii values for the Group 1A elements. Use the concept of effective nuclear charge to rationalize the trend in radii values for the Period 2 elements. You can see a short movie depicting what you should have concluded here.

Ionization Energy

Recall that the ionization energy (actually the first ionization energy) is the energy required to remove an electron from the outermost occupied orbital of a gaseous atom.

PROBLEMS

Plot the ionization energies for the first 86 elements of the periodic table versus atomic number.Display your plot as both a normal graph and as a bar graph. Explain the general trends that occur across a period and down a family using the concept of effective nuclear charge. In each period an anomaly to the general trend occurs with the Group 3A element and the Group 6A element. Explain this anomalous behavior. Using Slater's rules, calculate the Z* for Al, Al+, Al2+, and Al3+. Discuss the results relative to the expected ionization energies for these species. Calculate the Z* for the valence ns electron of lithium, sodium, and potassium using the assumption that σ equals the number of inner-shell electrons. Are your results consistent with the trends in ionization energy for these elements? Discuss why or why not.

Electron Affinity

Electron affinity is the change in energy that occurs when an electron is added to a neutral, gaseous atom.

PROBLEMS

Plot the electron affinity for Main Group elements ( Group 1A-8A) versus atomic number.Display your plot as both a normal graph and as a bar graph. Explain the general trends that occur across a period and down a family using the concept of effective nuclear charge. An anomaly to the general trend occurs in a period in going from the Group 1A element to the Group 2A element and in going from the Group 4A to Group 5A. An anomaly can also be found between rows 2 and 3 in going down a period. Give an explanation for each of these anomalies.

Electronegativity

Electronegativity is the ability of an atom in a molecule to draw electrons to itself.

PROBLEMS

Sketch a periodic table, indicating the trend of electronegativities from lowest to highest. Relate these trends to effective nuclear charge and atomic size.

Other Trends

There are many properties for which you can investigate periodic trends. There is a nice Excel spreadsheet that allows you to look at the trends for various properties in 3-D. Open the spreadsheet and view the plots.

If you have difficulty viewing the spreadsheet in conjunction with this webpage, try downloading it (right clicking on the link will allow you to save the file). This spreadsheet uses macros for accessing the plots. If you get a message that the macros have not downloaded because of the security setting , reset the security setting for Excel using the tools menu

A. DEVELOPMENT OF PERIODIC SYSTEM
Enterprises grouping based on common elements of nature done for these elements are easy to learn.

1. Triade Dobereiner
In 1829, Johan Wolfgang Dobereiner study the properties of some of the elements are already known at the time. Dobereiner see any resemblance between some elements of nature and the elements are grouped according to similarities in character. Apparently each group consists of three elements of the so-called triade. If the elements in the triade compiled and sorted based on the similarity of its atomic mass, then the second element is an average of the properties and atomic masses of the first and third elements.

2. Octave Theory Newland
In 1864, John Alexander Reina Newland compiled a list of more numerous elements. Newland arrangement shows that if the elements arranged by atomic mass increases, then the first element has similar properties to the eighth element, the second element is similar to the elements of nature ninth, and so on. The discovery was declared a law Newland Newland Octave.

At Newland Octave compiled list, the elements of noble gases (He, Ne, Ar, Kr, Xe, and Rn) has not been found. Noble gases are found by Rayleigh and Ramsay in 1894. Elements of the noble gases argon gas was first discovered. Newland Octave law only applies to elements with low atomic mass.

3. Mendeleev Periodic System
In 1869, began the periodic table arranged. The periodic table is the work of two scientists, Dmitri Ivanovich Mendeleev from Russia and Julius Lothar Meyer from Germany. They work independently and produce a similar table at the same time. Mendeleev presents its results to the Russian Chemical Society in early 1869, and a new periodic table Meyer appeared in December 1869.

Mendeleev who first suggested the periodic table, it is considered as the inventor of the periodic table is often referred to as short-periodic system of elements. Mendeleev developed the periodic table by atomic mass increases and similar properties. Mendeleev periodic system was first published in the scientific journal Annalen der Chemie in 1871.

The important thing there is in the Mendeleev periodic system are as follows:
a. two adjacent elements, has a relative atomic mass difference of at least two or one unit;
b. there are empty boxes for undiscovered elements, such as 44, 68, 72, and 100;
c. can predict the properties of elements that have not been known as ekasilikon;
d. measurement error can be corrected relative atomic mass of some elements, such as Cr = 52.0 not 43.3.

a. Excess Mendeleev periodic system
1) Chemical and physical properties in a single class of similar elements and changed regularly.
2) the highest valence of an element equal to their group number.
3) to predict the properties of elements yet to be discovered at the time and has had a vacant place.

b. Mendeleev periodic system deficiencies
1) The length of the period is not the same and why not described.
2) Some elements are not ordered by their atomic mass increases, for example: Te (128) before I (127).
3) the mass difference of successive elements are not always 2, but the range between 1 and 4 so difficult to predict the mass of an element that has not been known.
4) Valence more than one element of the unpredictable group.
5) anomaly (deviation) elemental hydrogen from other elements not described.

4. Modern Periodic System
In 1914, Henry G. J. Moseley found that the order of the elements in the periodic table according to the atomic number of the element increases. Moseley managed to find errors in the Mendeleev periodic table, that there is an inverse element is located. Placement tellurium and iodine is not in accordance with the increase of relative atomic mass, were consistent with the increase in atomic number.

Tellurium has the atomic number 52 and iodine has an atomic number of 53. The modern periodic system can be regarded as a refinement Mendeleev periodic system. The modern periodic system, also known as the long form of the periodic system, is based on the atomic number increases, and similar properties. In the modern periodic system of horizontal rows are called periods and upright rows are called groups.

The number of periods in the periodic system there are 7 and marked with numbers:
• Period 1 is referred to as the period is very short and contains 2 elements
• Period 2 is referred to as short a period and contains 8 elements
• Period 3 is referred to as short a period and contains 8 elements
• Period 4 called the period length and contains 18 elements
• Period 5 called the period length and contains 18 elements
• Period of six called the period is very long and contains 32 elements, in this period there are elements of the Lanthanide element number 58 to number 71 and placed on the bottom
• Period 7 is referred to as the period is not yet complete because it may increase the number of elements that occupy it again, until now contains 24 elements. In this period there is a row of elements called Actinides, ie elements numbered 90 to 103 and the number is placed at the bottom.

The number of groups in the periodic system are 8 and marked with Roman numerals. There are two major categories, namely class A (main group) and class B (class transitions). Group B is located between groups IIA and IIIA groups.

The names of the classes to the element of class A
• Group IA alkali group called
• Group IIA alkaline earth group called
• Type IIIA called boron golonga
• Group IVA called carbon groups
• type is called a class of nitrogen VA
• VIA Group called the oxygen group
• Group VIIA group called halogens
• Group VIIIA group called noble gases

In the period 6 class IIIB there are 14 elements that are very similar, the lanthanide elements. In the period 7 also apply the same thing and called the actinide elements. Both series of elements called transition elements inside.

The elements of the lanthanide and actinide belonged IIIB, included in the
group because it has very similar properties.

B. NATURE OF METAL
Properties owned by the elements very much. In this discussion, we will only discuss some properties of the elements. Based on the nature kelogamannya, generally elements are divided into three categories: metals, non-metallic elements, and the element metalloid (semi-metal).

Metals we have encountered many around us, for example iron, aluminum, copper, silver, gold, and others. In general, the metal has physical properties, such as:
1. good heat conductor;
2. good conductor of electricity;
3. shiny metal surfaces;
4. can be beaten into thin plates;
5. can stretch when pulled.

The ability of the metal to stretch when pulled is called ductility. The ability to conduct electricity and stretch metal used to make the wire or cable. The ability to change shape if forged metal called maleabilitas. The ability of a metal to change shape if forged dimanfaatka to make various kinds of goods, such as machetes, knives, hoes, and others. The above properties are not owned by the non-metal elements (non-metallic).

When viewed from the electron configuration, metallic elements tend to lose electrons (having a small ionization energy), while non-metallic elements tend to capture electrons (it has a large ionization energy).

Thus, we can see the tendency of metals in the periodic system properties, ie in a group from top to bottom is getting bigger and in a period from left to right are getting smaller. If we look at the periodic table of elements, metal elements berletak on the left, while the non-metallic elements located on the right (see the periodic table of elements).

On the periodic table, the boundary between metallic elements and nonmetallic often depicted with thick ribbed diagonal ladder. The elements in the border area have a dual nature. For example, the metal beryllium (Be) and aluminum (Al), the metals have some nonmetallic properties, and is called amphoteric elements. The metal which is next door (in the periodic table) are Boron (B) and Silicon (Si) is a non-metallic element, which owns several properties of metals and metalloid elements called

In the SparkNote on the Periodic table we discussed a number of simple periodic trends. In this section we will discuss a number of more complex trends, the understanding of which relies on knowledge of atomic structure.
Before getting into these trends, we should engage a quick review and establish some terminology. As seen in the previous section on the octet rule, atoms tend to lose or gain electrons in order to attain a full valence shell and the stability a full valence shell imparts. Because electrons are negatively charged, an atom becomes positively or negatively charged as it loses or gains an electron, respectively. Any atom or group of atoms with a net charge (whether positive or negative) is called an ion. A positively charged ion is a cation while a negatively charged ion is an anion.
Now we are ready to discuss the periodic trends of atomic size, ionization energy, electron affinity, and electronnegativity.

Atomic Size (Atomic Radius)

The atomic size of an atom, also called the atomic radius, refers to the distance between an atom's nucleus and its valence electrons. Remember, the closer an electron is to the nucleus, the lower its energy and the more tightly it is held.

Moving Across a Period

Moving from left to right across a period, the atomic radius decreases. The nucleus of the atom gains protons moving from left to right, increasing the positive charge of the nucleus and increasing the attractive force of the nucleus upon the electrons. True, electrons are also added as the elements move from left to right across a period, but these electrons reside in the same energy shell and do not offer increased shielding.

Moving Down a Group

The atomic radius increases moving down a group. Once again protons are added moving down a group, but so are new energy shells of electrons. The new energy shells provide shielding, allowing the valence electrons to experience only a minimal amount of the protons' positive charge.

Cations and Anions

Cations and anions do not actually represent a periodic trend in terms of atomic radius, but they do affect atomic radius, and so we will discuss them here.
A cation is positively charged, meaning that it is an atom that has lost an electron or electrons. The positive charge of the nucleus is thus distributed over a smaller number of electrons and electron-electron repulsion is decreased, meaning that the electrons are held more tightly and the atomic radius is smaller than in the normal neutral atom. Anions, conversely, are negatively charged ions: atoms that have gained electrons. In anions, electron-electron repulsion increases and the positive charge of the nucleus is distributed over a large number of electrons. Anions have a greater atomic radius than the neutral atom from which they derive.

Ionization Energy and Electron Affinity

The process of gaining or losing an electron requires energy. There are two common ways to measure this energy change: ionization energy and electron affinity.

Ionization Energy

The ionization energy is the energy it takes to fully remove an electron from the atom. When several electrons are removed from an atom, the energy that it takes to remove the first electron is called the first ionization energy, the energy it takes to remove the second electron is the second ionization energy, and so on. In general, the second ionization energy is greater than first ionization energy. This is because the first electron removed feels the effect of shielding by the second electron and is therefore less strongly attracted to the nucleus. If a particular ionization energy follows a previous electron loss that emptied a subshell, the next ionization energy will take a rather large leap, rather than follow its normal gently increasing trend. This fact helps to show that just as electrons are more stable when they have a full valence shell, they are also relatively more stable when they at least have a full subshell.
Ionization Energy Across a Period
Ionization energy predictably increases moving across the periodic table from left to right. Just as we described in the case of atomic size, moving from left to right, the number of protons increases. The electrons also increase in number, but without adding new shells or shielding. From left to right, the electrons therefore become more tightly held meaning it takes more energy to pry them loose. This fact gives a physical basis to the octet rule, which states that elements with few valence electrons (those on the left of the periodic table) readily give those electrons up in order to attain a full octet within their inner shells, while those with many valence electrons tend to gain electrons. The electrons on the left tend to lose electrons since their ionization energy is so low (it takes such little energy to remove an electron) while those on the right tend to gain electrons since their nucleus has a powerful positive force and their ionization energy is high. Note that ionization energy does show a sensitivity to the filling of subshells; in moving from group 12 to group 13 for example, after the d shell has been filled, ionization energy actually drops. In general, though, the trend is of increasing ionziation energy from left to right.

Ionization Energy Down a Group

Ionization energy decreases moving down a group for the same reason atomic size increases: electrons add new shells creating extra shielding that supersedes the addition of protons. The atomic radius increases, as does the energy of the valence electrons. This means it takes less energy to remove an electron, which is what ionization energy measures.

Electron Affinity

An atom's electron affinity is the energy change in an atom when that atom gains an electron. The sign of the electron affinity can be confusing. When an atom gains an electron and becomes more stable, its potential energy decreases: upon gaining an electron the atom gives off energy and the electron affinity is negative. When an atom becomes less stable upon gaining an electron, its potential energy increases, which implies that the atom gains energy as it acquires the electron. In such a case, the atom's electron affinity is positive. An atom with a negative electron affinity is far more likely to gain electrons.

Electron Affinities Across a Period

Electron affinities becoming increasingly negative from left to right. Just as in ionization energy, this trend conforms to and helps explain the octet rule. The octet rule states that atoms with close to full valence shells will tend to gain electrons. Such atoms are located on the right of the periodic table and have very negative electron affinities, meaning they give off a great deal of energy upon gaining an electron and become more stable. Be careful, though: the nobel gases, located in the extreme right hand column of the periodic table do not conform to this trend. Noble gases have full valence shells, are very stable, and do not want to add more electrons: noble gas electron affinities are positive. Similarly, atoms with full subshells also have more positive electron affinities (are less attractive of electrons) than the elements around them.

Electron Affinities Down a Group

Electron affinities change little moving down a group, though they do generally become slightly more positive (less attractive toward electrons). The biggest exception to this rule are the third period elements, which often have more negative electron affinities than the corresponding elements in the second period. For this reason, Chlorine, Cl, (group VIIa and period 3) has the most negative electron affinity.

Electronegativity

Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those two atoms are associated through a bond. Electronegativity is based on an atom's ionization energy and electron affinity. For that reason, electronegativity follows similar trends as its two constituent measures.
Electronegativity generally increases moving across a period and decreases moving down a group. Flourine (F), in group VIIa and period 2, is the most powerfully electronegative of the elements. Electronegativity plays a very large role in the processes of Chemical Bonding.