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Covalent bond

Ionic and Covalent Bonds

❶If the structure is not correct, calculate the formal charge on each of the ligand atoms. Iodine fluoride, IF, is another diatomic compound that should have some polarity.

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Usually, this is the part that's not the hydrogens. Use the steps below to name oxygen anions: Figure out how many anions are in the same series as yours. Anions are in the same series when you can add or subtract oxygens and keep the overall charge the same by changing the oxidation state of the cation the positively-charged part of the anion. A good guide to oxidation states is available here. The first is sulfite and the second is sulfate based on their number of oxygens.

If there are more than 2 anions in the same series, use the prefix "hypo" for the 1 with least oxygen and the prefix "per-" for the 1 with the most oxygen. These are named hypochlorite , chlorite , chlorate , and perchlorate respectively. Use the name of the oxygen anion to find the name of the acid. Now that you've named the oxygen anion, finding the name of the acid itself isn't difficult. Use the following naming rules: If the anion ends in -ate, the acid ends with -ric or -ic.

HClO 2 is chlorous acid from "chlorate," the name of the oxygen anion. If the anion ends in -ite, the acid ends with -ous. HClO 3 is chloric acid from "chlorite," the name of the oxygen anion.

Add hypo- or per- to the beginning if the anion had these prefixes. HClO 4 is perchlorous acid from "perchlorate," the name of the oxygen anion. Identify acid name exceptions. As with normal covalent compounds, certain acids get special names that don't conform to the normal naming rules. There's no easy way to learn these, but they're easy to memorize as you start to come across them. Below are a few examples: HCN is "hydrocyanic acid. That is a mistake in the article. It should be dihydrogen oxide, as there are two hydrogens and one oxygen, not the other way around.

Not Helpful 0 Helpful 5. Not Helpful 0 Helpful 0. Include your email address to get a message when this question is answered. Already answered Not a question Bad question Other. By using this service, some information may be shared with YouTube. Tips When a prefix ends with a vowel and the name of the atom starts with a vowel, you usually drop the atom's vowel to keep the word from having a strange pronunciation.

For example, it's "monoxide," not "monooxide. You can remember acid naming rules by using: My r ide has hydro lics a play off of hydraulics. Spr ite is delici ous. I ate something ic ky. So if you have an anion that ends with - ide , use hydro. An anion ending with - ite matches with the acid name - ous.

Chemical Nomenclature In other languages: Thanks to all authors for creating a page that has been read 52, times. Did this article help you? Cookies make wikiHow better. By continuing to use our site, you agree to our cookie policy. RS Raymond Smith May This is by far the most succinct guide I found.

AP Aggie Pearson Sep 14, I did not know how to name compounds in English and this helped a lot! FA Farhana Allani Jan 14, Thank you so much. HA Heba Ahmed Nov 15, I hope you guys do more of "how to do" stuff. I love you, wikiHow! This article helps me out a lot, thanks. AS Ahmed Saadeh Mar 21, A Anonymous Feb 20, Ilakkiya Kiritharan Oct 27, Some texts refer to a bond that is between covalent and ionic called a polar covalent bond.

There is a range of bond between purely ionic and purely covalent that depends upon the electronegativity of the atoms around that bond.

If there is a large difference in electronegativity, the bond has more ionic character. If the electronegativity of the atoms is more similar, the bond has more covalent character.

Lewis structures are an opportunity to better visualize the valence electrons of elements. In the Lewis model, an element symbol is inside the valence electrons of the s and p subshells of the outer ring. It is not very convenient to show the Lewis structures of the Transition Elements, the Lanthanides, or Actinides. The inert gases are shown having the element symbol inside four groups of two electrons symbolized as dots.

Two dots are above the symbol, two below, two on the right, and two on the left. The inert gases have a full shell of valence electrons, so all eight valence electrons appear. Halogens have one of the dots missing. It does not matter on which side of the symbol the dot is missing. Group 1 elements and hydrogen are shown with a single electron in the outer shell.

Group 2 elements are shown with two electrons in the outer shell, but those electrons are not on the same side. Group 3 elements have three dots representing electrons, but the electrons are spread around to one per position, as in Group 2 elements. Group 4 elements, carbon, silicon, etc.

Group 5 elements, nitrogen, phosphorus, etc. In only one position are there two electrons. So Group 5 elements such as nitrogen can either accept three electrons to become a triple negative ion or join in a covalent bond with three other items.

When all three of the unpaired electrons are involved with a covalent bond, there is yet another pair of electrons in the outside shell of Group 5 elements. Group 6 elements, oxygen, sulfur, etc. Group 7 elements have all of the eight outside electrons spaces filled except for one. The Lewis structure of a Group 7 element will have two dots in all four places around the element symbol except for one.

Let's start with two atoms of the same type sharing a pair of electrons. Chlorine atoms have seven electrons each and would be a lot more stable with eight electrons in the outer shell.

Single chlorine atoms just do not exist because they get together in pairs to share a pair of electrons. The shared pair of electrons make a bond between the atoms. In Lewis structures, the outside electrons are shown with dots and covalent bonds are shown by bars. This covalent bond between chlorine is one of the most covalent bonds known. A covalent bond is the sharing of a pair of electrons. The two atoms on either side of the bond are exactly the same, so the amount of "pull" of each atom on the electrons is the same, and the electrons are shared equally.

Next, let's consider a molecule in which the atoms bonded are not the same, but the bonds are balanced. Methane, CH 4 , is such a molecule. If there were just a carbon and a single hydrogen, the bond between them would not be perfectly covalent.

Hydrogen has a slightly lower electronegativity than carbon, so the electrons in a single H-C bond would, on average, be closer to the carbon than the hydrogen.

Carbon would be more negative. But the Lewis structure below shows that there are four hydrogens around a carbon atom, and that they are evenly separated. In the CH 4 molecule, the four hydrogen atoms exactly balance each other out. The Lewis structure of methane does not have any electrons left over.

The carbon began with four electrons and each hydrogen began with two electrons. Only the bars representing the shared pairs of electrons remain. The carbon now shares four pairs of electrons, so this satisfies the carbon's need for eight electrons in the outside shell. Each hydrogen has a single shared pair in the outside shell, but the outside shell of the hydrogen only has two electrons, so the hydrogen has a full outer shell also.

The Lewis structure as shown on the left is not the real thing. The hydrogens repel each other, so the shape of the methane molecule is really tetrahedral, but the effect is the same. The methane shape drawn in primitive 3-D to the right is a more accurate representation of the methane tetrahedral molecule. Carbons and hydrogens are nice and easy to write in Lewis structures, because each carbon must have four attachments to it and each hydrogen atom must have one and only one attachment to it.

When the bonds around a carbon atom go to four different atoms, the shape of the bonds around that carbon is roughly tetrahedral, depending upon what the materials are around the carbon. Carbons are also able to have more than one bond between the same two. Consider the series ethane C 2 H 6 , ethene C 2 H 4 , common name is ethylene , and ethyne C 2 H 2 , common name is acetylene.

In writing the Lewis structure of compounds, the bars representing bonds are preferred to the dots representing individual electrons. Every carbon has four bonds to it showing a pair of electrons to make eight electrons or four orbitals in the outer shell. Each hydrogen atom has one and only one bond to it for two electrons in the outer shell that occupies the only orbital that hydrogen has.

All of the outer shells are usually filled. While we are doing this, notice that the Lewis structure of a molecule will show the shape of the molecule. All of the bonds in ethane are roughly the tetrahedral angle, so all of the hydrogen atoms are equivalent. The bonds in acetylene make it a linear molecule.

The bonds in ethylene are somewhat trigonal around the carbons, and the carbons cannot twist around that bond as they can around a single bond, so that the molecule has a flat shape and the attachments to the carbons are not equivalent.

This is also true. You will see this in the study of organic chemistry. This type of difference between the positions of the hydrogen atoms is called cis - trans isomerism. The Lewis structure shows the shape of a molecule or polyatomic ion with the bonds to each atom drawn at 90 degrees right, left, up, and down from the atomic symbol and the non — bonded electrons as dots, usually in pairs, around the atomic symbol in the left, right, up, and down positions around the atom.

We could set up a group of general guidelines for the drawing of Lewis structures for simple molecules or polyatomic ions. Write all the atoms in the material in the form of the formula of the compound.

CO 2 can be an example. This is the proposed shape for the CO 2 molecule in the skeletal form. The proposed shape above has some problems with it. There are too many electrons assigned to the oxygen atoms and not enough to the carbon.

The way to express this idea is the formal charge. The formal charge is the number of electrons the atom brought to the structure minus the number of electrons shown in the proposed structure. The oxygen atoms both had six electrons in the valence shell because they are group VI A or group 16 atoms. They SHOW seven electrons in the proposed scheme, six dots and one electron from half the bond.

The carbon atom brought four electrons, being from group IV A or The formal charge of the carbon is plus two. The difference in formal charge indicates that there is a problem, but it also shows a likely way to balance things out. This process of writing Lewis structures is very limited to small molecules. There are many exceptions to the process, for instance, there are some compounds in which one atom has only three orbitals around it.

BF 3 , boron trifluoride is one in which the boron atom central is stuck with just three bonds to it. Some central atoms can have MORE than four orbitals around them. There is a phosphorus trichloride molecule PCl 3 that has the same shape as ammonia, but there is also a phosphorus penta chloride molecule PCl 5 that has five chlorine atoms attached to a central phosphorus. As you see, the scope of this tutorial goes only so far into the Lewis structure world.

The Lewis structures are usually good indicators of the actual shape of the molecule. We can tell that from the properties of the molecules. Rarely, but sometimes the best — looking Lewis structure is not the structure that predicts the properties of the material. In this case, the Lewis structure is wrong, and it probably makes some sense once the Lewis structure is written in the way that goes with the properties of the material. There is no issue of shape around the Group 1 elements.

There is only one attachment to them, so no angle is possible around them. But there are some molecular compounds with only two atoms, such as nitrogen monoxide, NO.

The only feature of this molecule is the bond between the nitrogen atom and the oxygen atom. The small difference in electronegativity between the oxygen and the nitrogen give the molecule a small dipole , a small separation of charge, so a small amount of polarity. Because there are an odd number of electrons in NO, this makes for an interesting Lewis structure.

Iodine fluoride, IF, is another diatomic compound that should have some polarity. Group 2 elements have two electrons in the outer shell.

Many of the compounds of Group 2 elements are ionic compounds, not really making an angle in a molecule. Molecules made with Group 2 elements that have two attached items to the Group 2 element have a linear shape, because the two attached materials will try to move as far from each other as possible. A linear shape means that a straight line could be made through all three atoms with the central element in the center. The shape of carbon dioxide is linear with the carbon in the center.

The idea is a disarmingly simple one. Electrons are all negatively charged, so they repel each other. If an atom has two electron groups around it, the electrons, and the atoms they are bonded to, are likely to be found as far as they can be from each other. Molecules with two electron groups attached to a central atom have a linear electron group shape and a linear molecular shape. Unless there is a large difference in electronegativity from one side to the other of a linear compound, there is no separation of charge and no polar character of the molecule.

Covalent compounds with boron are good examples of trigonal shaped molecules. The trigonal shape is a flat molecule with degree angles between the attached atoms.

Again using the example of a boron atom in the center, the attached elements move as far away from each other as they can, forming a trigonal shape, also called triangular , or trigonal planar to distinguish it from the trigonal pyramidal shape of compounds like ammonia.

BF3, boron trifluoride, is an example of a molecule with a trigonal planar shape. Each fluorine atom is attached to the central boron atom. There are three bonds to the boron, so the electron group shape is trigonal planar around boron. The molecular shap e is also trigonal planar in boron trifluoride because each electron group has a fluorine atom attached to it. But, what if the central atom has two other atoms and a lone pair of electrons attached to it?

Nitrogen oxychloride is an example of that. NOC l , is a molecule with nitrogen in the center See how to write Lewis structures above. When we go through the skeleton structure and distribute the electron dots, we find that there is a double bond between the nitrogen and the oxygen and a lone pair unshared pair of electrons on the nitrogen in addition to the single bond from the nitrogen to the chlorine.

There are three electron groups around the nitrogen, making the electron group shape more or less trigonal planar. But only two of those electron groups have an atom attached, so the molecular shape of nitrogen oxychloride is bent or angular. NOC l is not a balanced shape, so it is likely that there is some separation of charge within the molecule, making it a somewhat polar compound.

Group 4 elements are not in the center of a flat molecule when they have four equivalent attachments to them. As with two or three attachments, the attached items move as far as they can away from each other.

In the case of a central atom with four things attached to it, the greatest angle between the attached items does not produce a flat molecule.

If you were to cut off the vertical portion of a standard three-legged music stand so that it was the same length as the three legs, the angles among all four directions would be roughly equal. Try this with a gumdrop or a marshmallow. Stick four different colored toothpicks into the center at approximately the same angle.

If you have done it right, the general shape of the device will be the same no matter which one of the toothpicks is up. This shape is called tetrahedral. The shape of a tetrahedron appears with the attached atoms at the points of the figure and each triangle among any three of them makes a flat plane.

A tetrahedron is a type of regular pyramid with a triangular base. Carbon is a group four element. Methane, CH 4 , and carbon tetrachloride, CC l 4 , are good examples of tetrahedral shape.

If you draw the Lewis structures of these compounds, you will see that there are four bonds to the central carbon atom, but no other electrons on the central atom. They have four electron groups single bonds around the central atom, so they have a tetrahedral electron group shape.

Each bond to the central carbon has an atom attached, so they have a tetrahedral molecular shape. In both compounds, the four atoms attached to carbon are the same, so there is no separation of charge. All four atoms have the same electron pull in balanced directions, so these compounds are non — polar. Can a central carbon make molecules with other shapes around the central atom?

Yes, you remember carbon dioxide, where there are two double bonds around the carbon. The shape of around the acid carbons is trigonal planar because it has a double bond to it and only three electron groups, but the shape around the other carbons is tetrahedral.

In the Lewis structures the atoms are drawn at ninety degrees from each other, but the real shape around those carbons exists in three — space. Group 5 elements, for instance nitrogen or phosphorus, will become triple negative as they add three electrons in ionic reactions, but this is rare. Nitrides and phosphides do not survive in the presence of water. Covalent bonds with these elements do survive in water. The shape of the bonds and the lone pair of electrons around nitrogen and phosphorus is tetrahedral , just like the bonds around Group 4 elements.

The molecular shape is trigonal pyramidal. See the images below. The one on the left is a Lewis structure representation of an ammonia molecule. The one on the right is an attempt at showing the 3-D shape of the same ammonia molecule.

The color and the length of the bonds are only to show the shape better. Notice that the unshared pair lone pair of electrons actually repels MORE than the hydrogen atoms, so the angle between the hydrogen atoms is a little LESS than the tetrahedral angle of Group 6 elements, oxygen and sulfur, have six electrons in the valence shell.

The compounds they make usually have two pairs of unshared electrons. Just as in Group 5 elements, these pairs of unshared electrons serve as other attached atoms for the electron shape of the molecule. Group 6 elements make tetrahedral electron shapes, but now there are only two attached atoms. The angle between the hydrogens in water is about degrees. This peculiar shape is one of the things that makes water so special. Group 7 elements have only one chance of attachment, so there is not usually any shape around these atoms.

The alchemists of old had several other objectives aside from making gold. The thought of a fluid material that could dissolve anything, the universal solvent, was another alchemical project. No alchemist would say, though, what material would hold such a fluid. Surprisingly, the closest thing we have to a universal solvent is water. Water is not only a common material, but the range of materials it dissolves is enormous. The guiding principle for predicting which materials dissolve in which solvent is that 'like dissolves like.

Fluids with a separation of charge in the bonds will dissolve ionic materials. The bonds that hold hydrogen atoms to oxygen atoms are closer to covalent than ionic, but the bond does have a great deal of ionic character. Oxygen atoms are more electronegative than hydrogen atoms, so the electron pair is held closer to the oxygen atom.

Another way to look at it is that only a very small number of water molecules are ionized at any one time. Materials of a mildly covalent nature, such as small alcohols and sugars, are soluble in water due to the mostly covalent nature of the bonds in water. The shape of the water molecule is bent at about a degree angle due to the electron structure of oxygen. The two pairs of electrons that force the attached hydrogens into something close to a tetrahedral angle give the water molecule an unbalanced shape like a boomerang, with oxygen at the angle and the hydrogen atoms at the ends.

Since the oxygen atom pulls the electrons closer to it, the oxygen side of the molecule has a slight negative charge. Cations positive ions are attracted to the partial positive charge on the oxygen side of water molecules. Likewise, the hydrogen side of the molecule has a slight positive charge, attracting anions.

Polar materials such as salts, materials that have a separation of charge, dissolve in water due to the charge separation of water. The origin of the separation is called a dipole moment and the molecule itself can be called a dipole. The Lewis structure of water on the right above would almost tempt you to believe the molecular shape is linear. The actual shape is a little better shown as in the drawing on the right.

The oxygen has FOUR electron groups around it, making the electron group shape tetrahedral. The drawing shows a larger than ninety degree angle between the hydrogen atoms and the two pairs of unshared electrons lone pairs as having one pair coming out of the screen towards you and the other pair going into the screen. The oxygen has a larger electronegativity, so there is a larger concentration of electrons negative charge to the left of the molecule.

This dipole or separation of charge within the molecule makes water a polar solvent. It attracts positive ions to the oxygen side of the molecule and negative ions to the hydrogen side of the molecule.

Molecules or atoms that have no center of asymmetry are non-polar. Atoms such as the inert gases have no center of asymmetry. Molecules such as methane, CH 4 , are likewise totally symmetrical. Very small forces, called London forces, can be developed within such materials by the momentary asymmetries of the material and induction forces on neighboring materials.

These small forces account for the ability of non-polar particles to become liquids and solids. The larger the atom or molecule, the more potent the London forces, possibly due to the greater ability to separate charge within a larger particle.

Ionic Bonds

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Writing covalent compounds involves the use of Greek prefixes. They are listed in the table to the right. The prefixes are used to indicate the amount of .

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Describe how to write a formula for a covalent compound. Method to deduce the formulae of covalent compounds. Writing formulas for covalent compounds examples. 1. Atom T has a proton number of 6. Atom W has 10 neutrons and a nucleon number of Atoms T and W combine to form a compound. Determine the molecular formula of the .

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Writing & Naming Formulas of Ionic & Covalent Compounds © - Douglas Gilliland The Physical Science Series index 1. The 5 Steps for writing an ionic compound formula: (I)Write the symbols of the two elements. Polar covalent compounds have a partial charge at each end of the molecule. index Naming Covalent Compounds Solutions Write the formulas for the following covalent compounds: 1) antimony tribromide SbBr3 2) hexaboron silicide B6Si 3) chlorine dioxide ClO2 4) hydrogen iodide HI.

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Binary Covalent Compounds Between Two Nonmetals. Two nonmetals combine to form a covalent or molecular compound (i.e., one that is held together by covalent bonds which result from the sharing of electrons). In many cases, two elements can combine in several different ways to make completely different compounds. AP moiprods.tkr 8. Ionic Bonding – electrons are transferred between the atoms. Covalent Bonding – electrons are shared by the nuclei. Polar Covalent Bond –.