Na Valence Electrons

Posted on  by admin
  1. List Of Valence Electrons For Each Element
  2. Na Valence Electron Configuration
  3. Sodium Valence Electrons Ion

Learning Objectives

In order to write the Na electron configuration we first need to know the number of electrons for the Na atom (there are 11 electrons). When we write the configuration we'll put all 11 electrons in orbitals around the nucleus of the Sodium atom. In writing the electron configuration for sodium the first two electrons will go in the 1s orbital. Consider sodium: in its elemental form, it has one valence electron and is stable. It is rather reactive, however, and does not require a lot of energy to remove that electron to make the Na + ion. We could remove another electron by adding even more energy to the ion, to make the Na 2+ ion. In order to write the Na electron configuration we first need to know the number of electrons for the Na atom (there are 11 electrons). When we write the configuration we'll put all 11 electrons in orbitals around the nucleus of the Sodium atom. In writing the electron configuration for sodium the first two electrons will go in the 1s orbital.

  • To use Lewis electron dot symbols to predict the number of bonds an element will form.

Why are some substances chemically bonded molecules and others are an association of ions? The answer to this question depends upon the electronic structures of the atoms and nature of the chemical forces within the compounds. Although there are no sharply defined boundaries, chemical bonds are typically classified into three main types: ionic bonds, covalent bonds, and metallic bonds. In this chapter, each type of bond wil be discussed and the general properties found in typical substances in which the bond type occurs

  1. Ionic bonds results from electrostatic forces that exist between ions of opposite charge. These bonds typically involves a metal with a nonmetal
  2. Covalent bonds result from the sharing of electrons between two atoms. The bonds typically involves one nonmetallic element with another
  3. Metallic bonds These bonds are found in solid metals (copper, iron, aluminum) with each metal bonded to several neighboring groups and bonding electrons free to move throughout the 3-dimensional structure.

Each bond classification is discussed in detail in subsequent sections of the chapter. Let's look at the preferred arrangements of electrons in atoms when they form chemical compounds.

Lewis Symbols

At the beginning of the 20th century, the American chemist G. N. Lewis (1875–1946) devised a system of symbols—now called Lewis electron dot symbols (often shortened to Lewis dot symbols) that can be used for predicting the number of bonds formed by most elements in their compounds. Each Lewis dot symbol consists of the chemical symbol for an element surrounded by dots that represent its valence electrons.

Na valence electrons number

Lewis Dot symbols:

  • convenient representation of valence electrons
  • allows you to keep track of valence electrons during bond formation
  • consists of the chemical symbol for the element plus a dot for each valence electron

To write an element’s Lewis dot symbol, we place dots representing its valence electrons, one at a time, around the element’s chemical symbol. Up to four dots are placed above, below, to the left, and to the right of the symbol (in any order, as long as elements with four or fewer valence electrons have no more than one dot in each position). The next dots, for elements with more than four valence electrons, are again distributed one at a time, each paired with one of the first four. For example, the electron configuration for atomic sulfur is [Ne]3s23p4, thus there are six valence electrons. Its Lewis symbol would therefore be:

Fluorine, for example, with the electron configuration [He]2s22p5, has seven valence electrons, so its Lewis dot symbol is constructed as follows:

Lewis used the unpaired dots to predict the number of bonds that an element will form in a compound. Consider the symbol for nitrogen in Figure 8.1.2. The Lewis dot symbol explains why nitrogen, with three unpaired valence electrons, tends to form compounds in which it shares the unpaired electrons to form three bonds. Boron, which also has three unpaired valence electrons in its Lewis dot symbol, also tends to form compounds with three bonds, whereas carbon, with four unpaired valence electrons in its Lewis dot symbol, tends to share all of its unpaired valence electrons by forming compounds in which it has four bonds.

The Octet Rule

In 1904, Richard Abegg formulated what is now known as Abegg's rule, which states that the difference between the maximum positive and negative valences of an element is frequently eight. This rule was used later in 1916 when Gilbert N. Lewis formulated the 'octet rule' in his cubical atom theory. The octet rule refers to the tendency of atoms to prefer to have eight electrons in the valence shell. When atoms have fewer than eight electrons, they tend to react and form more stable compounds. Atoms will react to get in the most stable state possible. A complete octet is very stable because all orbitals will be full. Atoms with greater stability have less energy, so a reaction that increases the stability of the atoms will release energy in the form of heat or light ;reactions that decrease stability must absorb energy, getting colder.

When discussing the octet rule, we do not consider d or f electrons. Only the s and p electrons are involved in the octet rule, making it a useful rule for the main group elements (elements not in the transition metal or inner-transition metal blocks); an octet in these atoms corresponds to an electron configurations ending with s2p6.

Definition: Octet Rule

A stable arrangement is attended when the atom is surrounded by eight electrons. This octet can be made up by own electrons and some electrons which are shared. Thus, an atom continues to form bonds until an octet of electrons is made. This is known as octet rule by Lewis.

  1. Normally two electrons pairs up and forms a bond, e.g., (ce{H_2})
  2. For most atoms there will be a maximum of eight electrons in the valence shell (octet structure), e.g., (ce{CH_4})
The other tendency of atoms is to maintain a neutral charge. Only the noble gases (the elements on the right-most column of the periodic table) have zero charge with filled valence octets. All of the other elements have a charge when they have eight electrons all to themselves. The result of these two guiding principles is the explanation for much of the reactivity and bonding that is observed within atoms: atoms seek to share electrons in a way that minimizes charge while fulfilling an octet in the valence shell.

The noble gases rarely form compounds. They have the most stable configuration (full octet, no charge), so they have no reason to react and change their configuration. All other elements attempt to gain, lose, or share electrons to achieve a noble gas configuration.

Example (PageIndex{1}): Salt

The formula for table salt is NaCl. It is the result of Na+ ions and Cl- ions bonding together. If sodium metal and chlorine gas mix under the right conditions, they will form salt. The sodium loses an electron, and the chlorine gains that electron. In the process, a great amount of light and heat is released. The resulting salt is mostly unreactive — it is stable. It will not undergo any explosive reactions, unlike the sodium and chlorine that it is made of. Why?

Solution

Referring to the octet rule, atoms attempt to get a noble gas electron configuration, which is eight valence electrons. Sodium has one valence electron, so giving it up would result in the same electron configuration as neon. Chlorine has seven valence electrons, so if it takes one it will have eight (an octet). Chlorine has the electron configuration of argon when it gains an electron.

The octet rule could have been satisfied if chlorine gave up all seven of its valence electrons and sodium took them. In that case, both would have the electron configurations of noble gasses, with a full valence shell. However, their charges would be much higher. It would be Na7- and Cl7+, which is much less stable than Na+ and Cl-. Atoms are more stable when they have no charge, or a small charge.

Lewis dot symbols can also be used to represent the ions in ionic compounds. The reaction of cesium with fluorine, for example, to produce the ionic compound CsF can be written as follows:

List Of Valence Electrons For Each Element

No dots are shown on Cs+ in the product because cesium has lost its single valence electron to fluorine. The transfer of this electron produces the Cs+ ion, which has the valence electron configuration of Xe, and the F ion, which has a total of eight valence electrons (an octet) and the Ne electron configuration. This description is consistent with the statement that among the main group elements, ions in simple binary ionic compounds generally have the electron configurations of the nearest noble gas. The charge of each ion is written in the product, and the anion and its electrons are enclosed in brackets. This notation emphasizes that the ions are associated electrostatically; no electrons are shared between the two elements.

Atoms often gain, lose, or share electrons to achieve the same number of electrons as the noble gas closest to them in the periodic table.

As you might expect for such a qualitative approach to bonding, there are exceptions to the octet rule, which we describe elsewhere. These include molecules in which one or more atoms contain fewer or more than eight electrons.

Summary

Lewis dot symbols can be used to predict the number of bonds formed by most elements in their compounds. One convenient way to predict the number and basic arrangement of bonds in compounds is by using Lewis electron dot symbols, which consist of the chemical symbol for an element surrounded by dots that represent its valence electrons, grouped into pairs often placed above, below, and to the left and right of the symbol. The structures reflect the fact that the elements in period 2 and beyond tend to gain, lose, or share electrons to reach a total of eight valence electrons in their compounds, the so-called octet rule. Hydrogen, with only two valence electrons, does not obey the octet rule.

Contributors and Attributions

  • Mike Blaber (Florida State University)

  • National Programme on Technology Enhanced Learning (India)

Ionic Bonding

Soduim atom'>We have learned that atoms tend to react in ways that create a full valence shell, but what does this mean? Sodium (Na), for example, has one electron in its valence shell. This is an unstable state because that valence shell needs eight electrons to be full. In order to fill its valence shell, sodium has two options:

  1. Find a way to add seven electrons to its valence shell, or
  2. Give up that one electron so the next lower energy shell (already full) could become its new valence shell.

Which do you think is more easily accomplished?

Right, giving up one electron! Atoms like sodium, with only one or two electrons in a valence shell that needs eight electrons, are most likely to give up their valence electrons to achieve a stable state. All these atoms need is another atom that can attract their electrons!

An atom such as chlorine (Cl), that contains seven electrons in its valence shell, needs one more electron to have a full valence shell. If we use the same logic as we did for sodium, we should conclude that chlorine would rather gain one more electron than lose all seven of its valence electrons to achieve stability. Under the right conditions, atoms like chlorine will steal an electron from nearby atoms like sodium. This ability to pull electrons away from other atoms is termed electronegativity. Atoms with a valence shell that is almost full are more likely to be electronegative because they have greater reason to pull electrons towards them. Electronegative atoms are not negatively charged, but they are more likely to become negatively charged.

When an electron moves from one atom to another, both atoms become ions. An ion is any atom that has gained electrons to have a negative charge (anion) or lost electrons to have a positive charge (cation). An easy way to remember that a cation has a positive, or +, charge is to think of the letter t in the word cation as a + sign.

Once an atom becomes an ion, it has an electrical charge. Ions of opposite charge are attracted to one another, forming a chemical bond, an association formed by attraction between two atoms. This type of chemical bond is called an ionic bond because the bond formed between two ions of opposite charge. The sodium cation (Na+) and the chlorine anion (Cl-) are attracted to one another to form sodium chloride, or table salt.

Although ionic bonds are very strong, they can be relatively easily broken if another attractive ion (or polar molecule) comes around. An ionic bond is formed when two ions of opposite charge come together by attraction, NOT when an electron is transferred.

Think of ionic bond formation as the second step in a two step process:

  1. Two atoms each become ions. The atoms could have become ions in previous reactions with other atoms or the atoms may have reacted with each other, transferring the electron(s) from one to the other.
  2. The two ions of opposite charge “see” one another and become attracted enough to form the bond.

Ionic Bonding

This video visually illustrates how atoms form ionic bonds.

Create an Ionic Bond

Na Valence Electron Configuration

In this activity, you'll create cations and anions and watch the ionic bond form.

Sodium Valence Electrons Ion

Covalent Bonding

Carbon atom'>In ionic bonding, we looked at atoms with either one or two electrons in their valence shell and atoms that only needed one or two electrons to fill their valence shell. What happens when an atom, carbon (C) for example, has four valence electrons? Carbon would need to either lose four electrons or gain four electrons in order to have a full valence shell. Both of these situations would cause carbon to have a very strong charge, which would probably make it as unstable as having an incomplete valence shell! For atoms like carbon, there is another option: sharing.

When two atoms each need additional electrons to fill their valence shells, but neither is electronegative enough to steal electrons from the other, they can form another kind of chemical bond called a covalent bond. In covalent bonds, two atoms move close enough to share some electrons. The electrons from each atom shift to spend time moving around both atomic nuclei.

In the most common form of covalent bond, a single covalent bond, two electrons are shared, one from each atom’s valence shell. Double covalent bonds where four electrons are shared, and triple covalent bonds where six electrons are shared, are also commonly found in nature.

How do we know if and when covalent bonds will form? Atoms will form as many covalent bonds as it takes to fill their valence shell. This means that carbon, our previous example, will need to form four covalent bonds in order to fill its outermost shell. In each of the four bonds, carbon will contribute one electron and the other atom will contribute one electron, supplying carbon with eight electrons effectively orbiting its nucleus. Atoms like oxygen (O) will form two covalent bonds because they already have six valence electrons and only need two more electrons obtained by sharing. Put another way, oxygen shares two of its six valence electrons in covalent bonds while keeping four valence electrons for itself (4 unshared oxygen electrons + 2 shared oxygen electrons + 2 shared electrons from the other atoms = 8 total electrons).

In shell models, the shared electrons are shown within the overlapping region of the valence shells to represent the fact they are shared, but the electrons are actually moving around both nuclei and could be found anywhere around either nucleus at a given time.

For simplicity, we often draw covalent bonds as straight lines between atoms to represent the structural formula. Each line represents a single covalent bond (two shared electrons), so double lines represent a double covalent bond (four shared electrons).

Covalent bonds are usually found in atoms that have at least two, and usually less than seven, electrons in their outermost energy shell, but this is NOT a hard and fast rule. Hydrogen, for example, is a unique atom that bears closer examination. Because hydrogen only has one electron surrounding its nucleus, its valence shell is the first energy shell, which only needs two electrons to be full. Hydrogen tends to form covalent bonds because a single covalent bond will fill its shell. However, the one-proton nucleus is very weak and has trouble keeping the shared electrons around the hydrogen atom for very long.

Covalent Bonding

This video visually illustrates how atoms form covalent bonds.

Other Kinds of Bonding

When a covalent bond forms between atoms of similar electronegativity, the shared electrons tend to spend equal time around each nucleus. What happens if a bond is formed between atoms like oxygen, which is highly electronegative, and hydrogen, which is not? The oxygen atom tends to pull the shared electrons over to its side of the bond more often than the hydrogen atom does, resulting in polarity, or a partial separation of charge between atoms. The electrons don’t actually leave the less electronegative atom, but they do spend less time on that side. This results in two poles, one slightly positive and one slightly negative. This is called a polar bond, while covalent bonds where electrons are shared equally are called nonpolar.

Polar covalent bonds are the source of an additional kind of association called a hydrogen bond. In a hydrogen bond, the partially positive end of a polar covalent molecule is attracted to the partially negative end of another polar covalent molecule. For example, water is composed of two hydrogen atoms covalently bonded to a single oxygen atom. In a container full of water molecules, the hydrogen atoms of each water molecule are attracted to the oxygen atoms of other water molecules, forming hydrogen bonds between all of the molecules in the container of water. Hydrogen bonds are weak compared to covalent bonds, but they are strong enough to affect the behavior of the atoms involved. This leads to many important chemical properties in water and other molecules.