7.6 Molecular Structure and Polarity

• The lattice energy can be calculated from other values.

• Bond dissociation energies are used to calculate lattice energies for ionic compounds.
• The range of lattice energies is 600-4000 kJ/mol, whereas the dissociation bond energies are between 150 and 400 kJ/mol for single bonds.
• These are not directly comparable values.
• cations and anions pack together in an extended lattice, and lattice energies are associated with many interactions.
• The bond dissociation energy is associated with the interaction of two atoms.

• By the end of this section, you will be able to: Predict the structures of small molecule using valence shell electron pair repulsion (VSEPR) theory It is important to be able to describe bonds in terms of their distances, angles, and relative arrangements in space.
• Bond distances are measured in Angstroms or Picometers.

• The angles and bond distances are shown for the H2CO molecule.

• The model assumes that the arrangement of the electron pairs in the central atom will maximize the distance between them.
• The bonding pairs of electrons are located between the bonds of the atoms.
• The positions of the high electron density regions affect the repulsion of these electrons.

• The correct arrangement of atoms in a molecule is predicted by the theory.
• The theory only considers electron-pair repulsions.
• Nuclear-nuclear repulsions and nuclear-electron attractions are two of the interactions that are involved in the final arrangement of atoms.

• Predicting the structure of a BeF2 molecule is a simple example of VSEPR theory.
• There are only two electron pairs around the central beryllium atom in the Lewis structure of BeF2.
• With two bonds and no lone pairs of electrons on the central atom, the bonds are as far apart as possible, and the electrostatic repulsion between these regions of high electron density is reduced to a minimum when they are on opposite sides of the central atom.

• The bond angle is 180 degrees.

• The BeF2 molecule has two bonds on opposite sides of the Be atom in a linear structure.

• This and other electron-pair geometries minimize the repulsions among regions of high electron density.

• The basic electron-pair geometries maximize the space around any region of electron density.

• When there are no lone electron pairs around the central atom, the electron-pair geometries will be the same as the molecular structures, but they will be different.

• The methane molecule, CH4, which is the major component of natural gas, has four bonding pairs of electrons around the central carbon atom.
• The ammonia molecule, NH3 has four electron pairs associated with the nitrogen atom, and thus has a tetrahedral electron-pair geometry.
• The shape of the molecule is influenced by the lone pair, which is not included in the structure.

• The hydrogen atoms are arranged in a way that shows the structure of the methane molecule.
• Solid lines represent bonds in the plane of the page, solid wedges represent bonds coming up out of the plane, and dashed lines represent bonds going down into the plane, which are represented by the wedge and dash notation.

• The amount of space occupied by different kinds of electron pairs is predicted by the theory.

• The amount of space occupied by different regions of electrons is determined by the order of repulsions.
• The order of sizes from largest to smallest is: lone pair, triple bond, double bond, and single bond.
• There are two single bonds and one double bond in this molecule.
• The angle between the single bonds is slightly smaller than the double bonds, but the basic geometry is trigonal with 120deg bond angles.

• In the ammonia molecule, the three hydrogen atoms attached to the central nitrogen are arranged in a three-dimensional trigonal pyramid with the nitrogen atom at the apex and the three hydrogen atoms forming the base.
• The bond angles in a trigonal pyramid are based on the geometry of the electron pair.
• There are slight deviations from the ideal because lone pairs occupy larger regions of space.
• The H-N-H bond angles in NH3 are slightly smaller than in a regular tetrahedron because the lone pair-bonding pair repulsion is greater than the bonding pair-bonding pair repulsion.

• The structures are the same when there are no lone pairs.
• Modifications of the corresponding electron-pair geometry determine the structure of one or more lone pairs.

• The first three rows of the table are equivalent to the terminal atom locations.
• It doesn't matter if X is replaced with a lone pair or not because the molecule can be changed to convert positions.
• In a trigonal bipyramidal electron-pair geometry, lone pairs always occupy equatorial positions because these more spacious positions can more easily accommodate larger lone pairs.

• There are three possible arrangements for the three bonds and two lone pairs for the molecule, and the stable structure is the one that puts the lone pairs in equatorial locations.

• There are two lone electron pairs and four bonding regions in a central atom.

• If more than one arrangement of lone pairs and chemical bonds is possible, choose the one that will minimize repulsions, remembering that lone pairs occupy more space than multiple bonds.
• When every lone pair is in an equator position, repulsion is minimized.
• When the lone pairs are on opposite sides of the central atom, repulsion is minimized.

• The following examples show the use of VSEPR theory to predict the structure of a molecule with no lone pairs of electrons.
• The structure is the same as the electron pair geometry.

• There are two regions of high electron density around the carbon atom, and each double bond counts as one region.

• The electron-pair geometry is the same as the CO2 molecule.

• Three regions of high electron density give a trigonal geometry.
• The plane has 120 degree angles between the B-Cl bonds.
• BCl3 has a trigonal structure.

• Both the electron-pair geometry and the structure of BCl3 are trigonal.
• The Lewis structure shown above is not the correct bond shown by the VSEPR geometry.

• The electron-pair geometry is trigonal.
• All three C-O bonds are the same.
• Each bond counts as one region of electron density if it is single, double, or an average.

• Two of the top 50 chemicals produced in the United States contain the ion.

• The ion has a geometry and a structure.

• There is a molecule with trigonal bipyramidal.

• Any molecule with five electron pairs around the central atoms will be trigonal bipyramidal.
• It is a common example.

• Several examples show the effect of lone pairs of electrons.

• The electron-pair geometry is bent with an angle less than 109.5 degrees.
• The bond angle is 104.5%.

• When acids are dissolved in water, the hydronium ion forms.

• SF4 is used as a fluorinating agent in the preparation of compounds used as herbicides.

• The five regions are expected to adopt a trigonal bipyramidal electron-pair geometry.
• One of the positions is where the lone pair occupies.

• The trigonal bipyramidal is the electron-pair geometry.
• The structure is linear.

• xenon is the most reactive of the noble gases.

• The electron-pair geometry is adopted by the six regions.
• The five atoms are all in the same plane, so the lone pairs should be on opposite sides of the central atom.

• The central atom has three lone pairs and two bonds.

• The shape of the molecule is completely described by the structure of the molecule.
• The larger molecule has a chain of interior atoms that are connected by a local geometry.
• The way these local structures are oriented with respect to each other is an influence on the shape.

• We will only look at determining the local structures.

• The Lewis structure for H2NCH2CO2H is shown here.

• Consider each atom in its own way.

• The Lewis structure of alanine is shown here.

• You can build and name various molecules with the help of the lets.

• We can control whether bond angles and lone pairs are displayed by checking or unchecking the boxes on the right.
• We can hide or display the electron pair geometry by using the "Name" button at the bottom of the screen.

• Click on each bond type to add it to the central atom.
• You can examine the predicted structure of the molecule by rotating it.

• The structure is linear.

• A more complex molecule can be built.
• You can identify the electron-group geometry.
• Try to find a formula that matches the structure you have drawn.

• Answers will change.
• An atom with four single bonds, a double bond, and a lone pair has a square pyramidal molecular structure.
• The molecule XeOF4 adopts this structure.

• As discussed previously, polar covalent bonds connect two atoms with differing electronegativities, leaving one with a partial positive charge and the other with a partial negative charge, as the electrons are pulled toward the more positive atom.

• The arrows point along the bond from the less negative atom to the more positive atom.
• The partially positive end of the bond has a small plus sign drawn on it.
• The length of the arrow is determined by the electronegativity difference between the two atoms.

• A molecule may have a separation of charge depending on its structure.
• Adding the bond moments in three-dimensional space helps us determine the dipole moment.

• There is only one bond for diatomic molecules.

• The dipole moment of Homonuclear diatomic molecules is zero.
• There is a small dipole moment for CO.
• There is a larger 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609-

• The geometry must be taken into account when a molecule has more than one bond.
• If the bonds in a molecule are arranged in a way that their bond moments don't stay the same, then the molecule is nonpolar.
• The bonds in CO2 are polar, but the molecule as a whole is non polar.
• The CO2 molecule is linear with polar C-O bonds on the opposite side of the carbon atom.
• The bond moments are pointed in different directions.
• The Lewis structure shows that there are two bonds to a central atom, and the electronegativity difference shows that each of these bonds has a nonzero bond moment.
• This OpenStax book is available for free because of the lone pairs on O, and the two bond moments do not cancel.
• Water has a net dipole moment and is a polar molecule.

• The individual bond dipole moments of a molecule are dependent on how they are arranged.

• One of the oxygen atoms has been replaced by a sulfur atom in the OCS molecule.
• We draw the structure of the molecule to determine if it is polar.
• The C-O bond is considerably polar.
• The C-S bond is slightly polar because S is slightly more negative than C. The negative end of the molecule is the oxygen end.

• Another example of a polar molecule is chloromethane.
• The C-H bonds have a larger bond moment than the C-Cl bonds, and the bond moments do not completely cancel each other.

• In which all the polar bonds are identical, the molecules are nonpolar.
• The bonds in these molecules are arranged in such a way that their dipoles don't work.
• The dipoles will always cancel, even if a molecule contains identical bonds.
• The bonds and lone pairs on the central atoms are not canceled.
• H2S and NH3 are examples.

• The sum of the vectors of each bond does not cancel if you have a structure like that.

• When placed in an electric field with a positive end of a molecule and a negative end of a molecule, the molecule tends to align.
• We can use an object that is charged with an electric current to attract polar and non polar molecules.
• Both polar and non polar solvent are better at dissolving polar substances.

• The "Three Atoms" tab at the top of the book is available for free.
• This should show a molecule with three electronegativity adjustors.
• There are partial charges at the right and bond moments at the left.

• If A and C are very negative and B is in the middle of the range, use the electronegativity controls to determine how the molecule will look.

• The largest possible bond dipoles will be given by partial charges.

• The largest charges will be the largest bond moments.
• The two solutions show how different the electrons are in the bond.
• When the electronegativity difference is greatest, the bond moments will be maximized.
• The controls for A and C should be set to different extremes.
• The direction of the bond moment will not change based on whether B is the most positive or negative.

• Stable electron configurations are what atoms gain or lose electrons to form.
• The electronic structures of the cations formed by the representative metals have either a noble gas configuration or a completely filled electron shell.
• The charges of anions formed by nonmetals can be determined because they form when nonmetal atoms gain enough electrons.

• When electrons are shared between atoms, they 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- The electrons are shared equally in pure covalent bonds.
• One atom exerts a stronger force on the electrons than the other in polar covalent bonds.
• The ability of an atom to attract a pair of electrons is called electronegativity.
• polar bonds are determined by the difference in electronegativity between two atoms.
• There is no difference in electronegativity between the two atoms in a diatomic molecule.
• The bonding between metals and nonmetals is characterized as ionic when the electronegativity difference is large.

• Lewis symbols can be drawn for atoms and monatomic ion.
• Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the electrons are located in a Lewis structure.
• The octet rule states that every atom is surrounded by eight electrons.

• If half of the electrons are assigned to each atom, formal charges can be assigned to each atom.
• The most appropriate Lewis structure is determined by these hypothetical formal charges.
• The charges should be as close to zero as possible.
• There are two Lewis structures with the same arrangement of atoms but different distributions of electrons.
• The average of the distribution of electrons indicated by the individual Lewis structures is the resonance hybrid.

• The amount of energy required to break a bond in a mole is known as the bond energy dissociation.
• Multiple bonds are stronger than single bonds.
• The enthalpy of a reaction can be estimated using the energy input required to break bonds and the energy released when new bonds are formed.
• The lattice energy is the amount of energy needed to separate one mole of a compound from its gas phase ion.
• The OpenStax book is free and can be found at http://cnx.org/content/col11760/1.9 of the energetic steps involved in converting elements into an ionic compound.

• The three-dimensional arrangement of atoms is predicted by the theory.
• valence electrons will assume an electron-pair geometry that reduces repulsions between areas of high electron density When there are no lone electron pairs around the central atom, the placement of atoms in a molecule is equivalent to electron-pair geometry.
• A separation of charge is measured.
• The bond dipole moment is determined by the difference in electronegativity between the two atoms.
• The overall dipole moment of a molecule is determined by how the bonds are arranged in the structure.
• Nonpolar molecule interact with electric fields whereas polar molecule do not.

• A list of six ionic compounds can be prepared from the labels of several commercial products.
• Write the formula for each compound.

• The monatomic ion formed from the following list of elements is found in seawater.

• M and X represent elements in the third period of the periodic table.

• The Lewis structure for the diatomic molecule P2 is needed.

• The bonds in solid PbCl2 are ionic and the bonds in a HCl molecule are covalent.

• Some race cars use Methanol as their fuel.
• In Brazil, C2H5OH is used as motor fuel.
• CO2 and H2O are produced when they burn.
• Lewis structures are used to write the chemical equations for these reactions.

• Many planets in our solar system contain organic chemicals such as methane, ethane, propyne, and diacetylene.

• Fire extinguishers used to be used for electrical fires.
• Due to the formation of the toxic gas phosgene, it is no longer used for this purpose.

• The following electron configurations have the same atoms.

• The arrangement of atoms is given here.
• Add multiple bonds and lone pairs to complete the Lewis structures.
• Don't add any more atoms.

• A compound with a mass of about 28 g/mol contains 85.7% carbon and 14.3% hydrogen.

• A compound with a mass of about 42 g/mol contains 85.7% carbon and 14.3% hydrogen.

• Two arrangements of atoms are possible for a compound with a mass of 45 g/mol that contains 52.2% C, 13.1% H, and 34.7% O by mass.

• Write resonance forms that show the distribution of electrons.

• Write resonance forms that show the distribution of electrons.

• O3 is the component of the upper atmosphere that protects the Earth from ultraviolet radiation.

• bacon and other meats have been preserved with the ionic compound sodium nitrite.

• Explain why acetic acid, CH3CO2H, contains two different types of carbon-oxygen bonds, whereas the acetate ion only contains one type of carbonoxygen bond.

• The Lewis structures should be written for the following.

• Toothpastes containing hydrogen carbonate and hydrogen peroxide are widely used.
• Lewis structures for the hydrogen carbonate ion and hydrogen peroxide molecule can be written.

• Draw the structure of H3NO and assign formal charges.

• There is a series of fluorides formed by Iodine.
• Determine the formal charge of the atoms in this compound by writing the Lewis structure and chemical formula of the compound with a 70 g/mol molar mass.

• The industrial chemical sphuric acid is the most produced in the world.
• 90 billion pounds are produced in the United States alone.
• The Lewis structure for sulfuric acid has two oxygen atoms and two OH groups.

• The structure with the stronger bonds is usually the more stable form when a molecule can form two different structures.

• The bond energy of the carbon-sulfur double bond is calculated using the standard enthalpy of formation data.

• The first energy of Al is 578 kJ/mol and the first energy of Mg is 738 kJ/mol.

• The Educational Testing Service gave permission for this question to be taken from the chemistry exam.

• The lattice energy of KF is 794 kJ/mol.
• The NaF has the same structure as KF.

• If there is one or more lone pairs of electrons on M, the molecule is polar.
• There are two exceptions to this rule.

• There is a molecule with a dipole moment.

• There is a molecule with a dipole moment.

• There are three possible structures for PCl2F3.
• Draw them and discuss how they can be distinguished.

• A molecule with the formulaAB2 could have one of three different shapes.
• Pick out the three different shapes that this molecule might have.
• For each shape, give an example of a molecule or ion.

• A molecule with the formulaAB3 could have one of three different shapes.
• Pick out the three different shapes that this molecule might have.
• An example of a molecule with each shape is given.

• A compound with a molar mass of about 42 g/mol has 85.7% carbon and 14.3% hydrogen.

• Determine what values are needed to switch the dipole so that it points toward A.

• The following exercises can be performed using the simulation.
• To see certain dipoles, you may need to rotate the molecule in three dimensions.

• If you use the dipoles, you can predict whether N or H is more negative.

• You can check the box to see if you have a hypothesis.

• You can use the Molecule Shape Simulator to build a molecule.
• To add a double bond, click on the central atom.
• Add one bond and one pair.
• The complete geometry can be observed by rotating the molecule.
• Predict the bond angle by naming the electron group geometry.
• To check your answers, click the check boxes at the bottom and right of the simulation.