Chapter 3 - Chemical Bonds
Terms in this set (73)
- An attractive force that holds together the atoms, ions, or groups of atoms in a molecule or compound.
- results in a lower energy than the total energy of the separate atoms.
- if the lowest energy can be achieved by the complete transfer of one or more electrons from each atom of one element to those of another, then ions form and the compound is held together by the electrostatic attraction between them. This is called an ionic bond.
- if the lowest energy can be achieved by sharing electrons, then the atoms link through a covalent bond and discrete molecules are formed.
polar covalent bond
A covalent bond between atoms that differ in electronegativity. The shared electrons are pulled closer to the more electronegative atom, making it slightly negative and the other atom slightly positive. Property applies to nonmetal elements.
nonpolar covalent bond
a covalent bond in which the electrons are shared equally by the two atoms. Property applies to nonmetal elements.
The electrons in the outermost shell (main energy level) of an atom; these are the electrons involved in forming bonds.
A model used to calculate theoretical lattice enthalpies which assumes that the only bonding in an ionic solid is the electrostatic forces between the ions. These forces depend on the ionic radii and charges.
a kind of crystalline solid; aggregates of positively and negatively charged ions; no discrete ions; high melting points, high boiling points, poor electrical conductivity in the solid phase (due to strong electrostatic interactions which also cause ions to be relatively immobile).
-solids that consist of atoms, molecules or ions stacked together in a regular pattern.
atoms react by gaining or losing electrons so as to acquire the stable electron structure of a noble gas, usually eight valence electrons
duplet rule states that when an element gets 2 electrons in its valence or last shell it has achieved a stable electronic configuration . Hydrogen seems to be the only element obeying it, as well as Lithium, Berylllium.
- When atoms of metals on the left of the p-block in Periods 2 and 3 lose their valence electrons, they form ions with the electron configuration of the preceding noble gas.
- many metallic elements, such as those in the p- and d- blocks, have atoms that can lose a variable number of electrons.
- the ability of an element to form ions with different charges
- whenever two electrons are paired together in an orbital, or their total spin is 0, they are diamagnetic electrons. Atoms with all diamagnetic electrons are called diamagnetic atoms. A paramagnetic electron is an unpaired electron.
- nonmetals rarely lose electrons in chemical reactions because their ionization energies are too high. However, a nonmetal atom can acquire enough electrons to complete its valence shell and form an anion with an octet corresponding to the configuration of the next noble gas.
- To predict the electron configuration of a monatomic cation, remove the outermost electrons in the order np, ns, and (n-1)d; for a monatomic anion, add electrons until the next noble-gas configuration has been reached. The transfer of electrons results in the formation of an octet (or duplet) of electrons in the valence shell on each of the atoms: metal atoms achieve an octet (or duplet) by electron loss and nonmetal atoms achieve it by electron gain.
Consists of a chemical symbol for the element plus a dot for each valence electron. The valence electrons in any representative element is the same as the group number of the element. For example, Oxygen and Sulfur have 6 dots.
Writing ionic compounds with Lewis Symbols
1. Represent the cation by removing the appropriate number of dots from the symbol for the metal atom
2. Represent the anion by transferring those dots to the Lewis symbol for the nonmetal atom to complete its valence shell
3. If necessary, adjust the numbers of atoms of each kind so that all the dots removed from the metal atom symbols are accommodated by the nonmetal atom symbols
4. Write the charge of each ion as a superscript in the normal way
- Formulas of compounds consisting of the monatomic ions of main-group elements can be predicted by assuming that cations have lost all their valence electrons and anions have gained electrons in the their valence shells until each ion has an octet of electrons, or a duplet in the case of H, Li and Be.
- effective nuclear charge
- The effective nuclear charge is the net charge an electron experiences in an atom with multiple electrons.
Higher energy electrons can have other lower energy electrons between the electron and the nucleus, effectively lowering the positive charge experienced by the high energy electron.
the energy released when one mole of an ionic crystalline compound is formed from gaseous ions.
- the difference in energy between the ions of a compound widely separated as a gas and packed together in a solid -- usually very large.
- typically, only metallic elements have ionization energies that are low enough for the formation of ionic bonds to be energetically feasible.
- the energy required for the formation of ionic bonds is supplied largely by the attraction between oppositely charged ions.
- an ionic solid is not held together by bonds between specific pair sof ions -- an ionic bond is a "global" interaction characteristic of the entire crystal -- results in a net lowering of energy of the entire crystal relative of the widely separated neutral atoms.
- ionic solids have high melting points and are brittle
- high lattice energy indicates that the ions interact strongly with one an other to give a tightly bonded solid
- each ion in a solid experiences attractions from all other oppositely charged ions and repulsions form all other like-charged ions. The total potential energy is the sum of all these contributions.
- repulsions and attractions become progressively weaker as the distance from the central ion increases, but because the nearest neighbors of an ion give rise to a strong attraction, the net outcome of all these contributions is a lowering of energy.
- Negative potential energy = attraction
- when the potential energy is negative, there is a net lowering of energy, which means that the attraction between opposite charges overcomes the repulsion between like charges. The potential energy is strongly negative when the ions are highly charged and the separation between them is small, which is the case when the ions themselves are small.
- the factor A is a numerical coefficient called the Madelung constant; its value depends on how the ions are arranged about one another,
- ionic solids typically have high melting points and are brittle. The lattice energy of an ionic solid is large when the ions are small and highly charged.
- in covalent bond formation, atoms go as far as possible toward completing their octets by sharing electron pairs.
- the valence of an element is the number of bonds that its atoms can form
- pairs of valence electrons that do not take part in bonding
- the Lewis structure of a molecule shows atoms by their chemical symbols, covalent bonds by lines, and lone pairs by pairs of dots.
- nonmetal atoms share electrons until each has completed its octet (or duplet); a Lewis structure shows the arrangement of electrons as lines (bonding pairs) and dots (lone pairs).
- double and triple bonds are collectively called multiple bonds
- is the number of bonds that link a specific pair of atoms
- When making Lewis structures, make sure to choose as the central atom the element with the lowest ionization energy. This arrangement often results in the lowest energy because an atom in the central position shares more of its electrons than does a terminal atom. Atoms with higher ionization energies are more reluctant to share and are more likely to hold on to their electrons as lone pairs.
- another rule of thumb for predicting the structure of a molecule is to arrange the atoms symmetrically around the central atom.
- another clue for writing the correct arrangement of atoms is that, in simple chemical formulas, the central atom is often written first, followed by the atoms attached to it.
- for a cation, we subtract one dot for each positive charge. For an anion, we add one dot for each negative charge. The cation and the anion must be treated separately: they are individual ions and are not linked by shared pairs.
- the Lewis structure of a polyatomic species is obtained when all the valence electrons are used to complete the octets (or duplets) of the atoms present by forming single or multiple bonds and leaving some electrons as lone pairs.
- a blending of all three Lewis structures, with each bond intermediate in properties between a single and double bond. This blending of structures, which is called resonance, is depicted by double-headed arrows. The blended structure is a resonance hybrid of the contributing Lewis structures.
- Resonance is a mental exercise and method within the Valence Bond Theory of bonding that describes the delocalization of electrons within molecules. It compares and contrasts two or more possible Lewis structures that can represent a particular molecule. Resonance structures are used when one Lewis structure for a single molecule cannot fully describe the bonding that takes place between neighboring atoms relative to the empirical data for the actual bond lengths between those atoms. The net sum of valid resonance structures is defined as a resonance hybrid, which represents the overall delocalization of electrons within the molecule. A molecule that has several resonance structures is more stable than one with fewer. Some resonance structures are more favorable than others.
- electrons that are shown in different positions in a set of resonance structures are said to be delocalized. Delocalization means that a shared electron pair is distributed over several pairs of atoms and cannot be identified with just one pair of atoms. As well as delocalizing electrons over the atoms, resonance also lowers the energy below that of any single contributing structure and helps to stabilize the molecule. This lowering of energy occurs for quantum mechanical reasons. Broadly speaking, the wave function that describes the resonance structure is a more accurate description of the electronic structure of the molecule than the wave function for any single structure alone, and the more accurate the wave function, the lower the corresponding energy.
rules for writing appropriate resonance structures
1. In each contributing structure, the nuclei are in the same positions; only the locations of lone pairs and bonding pairs are changed.
2. Structures with the same energy (so called "equivalent structures") contribute equally to the resonance
3. Low-energy structures contribute more to the resonance mixture than high energy structures
- Kekule determined the structure of benzene. It consisted of conjugated double and single bonds in a circle of six carbons
resonance explains the different properties of benzene ...
1. Reactivity: Benzene does not undergo reactions typical of compounds with double bonds
2. Bond lengths: All the carbon-carbon bonds in benzene are the same length
3. Structural evidence: One 1,2 dichlorobenzene (in which chlorine atoms are attached to two adjacent carbon atoms) exists
- an important consequence of resonance is that is stabilizes a molecule by lowering its total energy. This stabilization makes benzene less reactive than expected for a molecules with three carbon-carbon double bonds
- Resonance is a blending of structures with the same arrangement of atoms but different arrangement of electrons. It spreads multiple-bond character over a molecule and results in a lower energy.
- one way to decide which structures are likely to make the major contribution is to compare the number of valence electrons distributed around each atom in a structure with the number of valence electrons on the free atoms. The smaller these differences for a structure, the lower is its energy and the greater is its contribution to a resonance hybrid.
- a measure of the redistribution of electrons is the formal charge on an atom in a given Lewis structure - the charge it would have if the bonding were perfectly covalent in the sense that the atom had exactly a half-share in the bonding electrons
- FC is assigned by establishing the "ownership" of the valence electrons of each atom in a molecule and comparing that ownership with the number of valence electrons on the free atom. An atom own one electron of each bonding pair attached to it and owns its lone pairs completely. The most plausible Lewis structure will be the one in which the formal charges of the atoms are lowest
- a Lewis structure in which the formal charges of the individual atoms are closest to zero typically represents the lowest-energy arrangement of the atoms and electrons
- Formal charge exaggerates the covalent character of bonds by assuming that the electrons are shared equally
- oxidation number exaggerates the ionic character of bonds. It represents the atoms as ions, and all the electrons in a bond are assigned to the atoms with the lower ionization energy (i.e. the more electronegative atom).
- the formal charge gives an indication of the extent to which atoms have gained or lost electrons in the process of covalent bond formation; atom arrangements and Lewis structures with the lowest formal charges are likely to have the lowest energy.
- some species have an odd number of valence electrons, and so at least one of their atoms cannot have an octet. Species having electrons with unpaired spins are called radicals. They are generally highly reactive.
- radicals are responsible for the rancidity of foods and the degradation of plastics in sunlight. Damage from radicals can be delayed by an additive called an antioxidant, which reacts rapidly with radicals before the radicals have a chance to do their damage.
- a molecule with two unpaired electrons. The unpaired electrons are usually on different atoms
- relative orientations of the two electron spins are random
- a radical is a species with an unpaired electron; a biradical has two unpaired electrons on either the same or different atoms.
expanded valence shell
Many molecules (and ions) have more than eight valance electrons around the central atom. That atom expands its valance shell to form more bonds, which release energy. These occur only with nonmetals from period 3 or higher because they have d orbitals available. The central atom may be bonded to more than four atoms or to four or fewer.
- the electrons in such an expanded valence shell may be present as lone pairs or may be used by the central atom to form additional bonds
- because the additional electrons must be accommodated in valence orbitals, only nonmetal atoms in Period 3 or later periods can expand their valence shell. Another factor - possibly the main factor - in determining whether more atoms than allowed by the octet rule can bond to the central atom is the size of that atom
- a compound that contains an atom with more atoms attached to it than is permitted by the octet rule
- elements that can expand their valence shell commonly show variable covalence, the ability to form different numbers of covalent bonds. Elements that have variable covalence can form one number of bonds in some compounds and a different number in others
- elements in period 3 and higher have empty 3d orbitals available, and so can expand their valence shells to accept additional electrons.
- when different resonance structures are possible, some giving the central atom in a compound an octet and some an expanded valence shell, the dominant resonance structure is likely to be the one with the lowest formal charges. However, there are many exceptions, and the selection of the best structure often depends on a careful analysis of experimental data.
- expansion of the valence shell to more than eight electrons occurs in elements of Period 3 and later periods. These elements can exhibit variable covalence and be hypervalent. Formal charge helps to identify the dominant resonance structure.
- an unusual feature of the Lewis structure of the colorless gas boron trifluoride is that the boron atom has an incomplete octet: its valence shell consists of only six electrons
- because fluorine has a high ionization energy, it is unlikely to exist with a positive formal charge. Experimental evidence, such as the short B-F bond lengths, suggests that the true structure of BF3 is a resonance hybrid of both types of Lewis structures, with the singly bonded structure making the major contribution.
Coordinate covalent bond
- the boron atom in BF3 can complete its octet if an additional atom of ion with a lone pair of electrons forms a bond by providing both electrons. A bond in which both electrons come form one of the atoms is called a coordinate covalent bond.
linked pairs of molecules.
- example is Al2Cl6. This molecule exists because a Cl atom in one AlCl3 molecule uses one of its lone pairs to form a coordinate covalent bond to the Al atom in a neighboring AlCl3 molecule. This arrangement can occur in aluminum chloride but not boron trichloride because the atomic radius of Al is bigger than that of B.
- compounds of Boron and Aluminum may have unusual Lewis structures in which Boron and Aluminum have incomplete octets or halogen atoms act as bridges.
- to describe bonds between nonmetals, covalent bonding is a good model. When a metal and nonmetal are present in a simple compound, ionic bonding is a good model
- the charges on the atoms in HCL are called partial charges. We show the partial charges on the atoms. A bond in which ionic contributions to the resonance result in partial charges are called polar covalent bonds. All bonds between atoms of different atoms are polar to some extent. The bonds in homonuclear (same-element) diatomic molecules and ions are nonpolar. The two atoms in a polar covalent bond form an electric dipole, a partial positive charge next to an equal but opposite partial negative charge.
- a dipole is represented by an arrow that points toward the positive partial charge
- the size of an electric dipole - which is a measure of the magnitude of the partial charges and their separation - is reported as the electric dipole moment.
- the SI unit of a dipole moment is 1 Cm (coulomb * meter)
- because ionization energies and electron affinities are highest at the top right of the periodic table, it is not surprising to find that nitrogen, oxygen, bromine, chlorine, and fluorine are the elements with the highest electronegativities.
- as the difference in electronegativities increases, so do the partial charges. If the difference in electronegativities is large, then one atom can acquire the lion's share of the electron pair, and the corresponding ionic structure make a large contribution to the resonance. Because it has largely robbed the other atom of its share of the electrons, the highly electronegative element resembles an anion and the other atom resembles a cation. We say that such a bond has considerable ionic character.
- an electronegativity difference of about 2 means that the bond has so much ionic character that it is best regarded as ionic.
- electronegativity is a measure of the pulling power of an atom on the electrons in a bond. A polar covalent bond is a bond between two atoms with partial electric charges arising from their difference in electronegativity. The presence of partial charges gives rise to an electric dipole moment.
- as the cation's positive charge pulls on the anion's electrons, the spherical electron cloud of the anion becomes distorted in the direction of the cation.
- atoms and ions with electron clouds that readily undergo large distortions are said to be highly polarizable
- atoms and ions that can cause large distortions are said to have high polarizing power.
- compounds composed of a small, highly charged cation and a large, polarizable anion tend to have bonds with considerable covalent character. Cations become smaller, more highly charged, and hence more strongly polarizing from left to right across a period. On the other hand, cations become larger and hence less strongly polarizing down a group.
- diagonal neighbors have similar polarizing properties
- ionic bonds acquire more covalent character as the distortion of the electron cloud on the anion increases
- compounds composed of highly polarizing cations and highly polarizable anions have significant covalent character in their bonding
- the strength of a chemical bond is measured by its dissociation energy, D, the energy required to separate the bonded atoms.
- the bond breaking is hemolytic, which means that each atom retains one of the electrons from the bond.
- a high dissociation energy indicates a deep potential energy well and therefore a strong bond that requires a lot of energy to break.
- the strength of a bond between two atoms is measured by its dissociation energy: the greater the dissociation energy, the stronger the bond
- a multiple bond is almost always stronger than a single bond, because more electrons bind the multiply bonded atoms. A triple bond between two atoms is always stronger than a double bond between the same two atoms, and a double bond is always stronger than a single bond between the same two atoms. However, a double bond between two carbon atoms is not twice as stronger as a single bond, and a triple bond is a lot less than three times as strong.
- the origin of these differences is in part the repulsions between the electron pairs in a multiple bond, so each pair is not quite as effective at bonding as a pair of electrons in a single bond. Another contributing factor is that the electron in double and c=triple bonds are not as concentrated between the two atoms as they are in a single bond.
- resonance spreads multiple-bond character over the bonds between atoms; as a result, what were single bonds are strengthened and what were double bonds are weakened. The net effect overall is a stabilization of the molecule.
- the presence of lone pairs may influence the strengths of bonds. Lone pairs repel each other; an, if they are on neighboring atoms, the repulsion can weaken the bond.
- trends in bond strength correlate with trends in atomic radii. If the nuclei of the bonded atoms cannot get very close to the electron pair lying between them, the two atoms will be only weakly bonded together.
- closely related to the strength of a bond is its stiffness (its resistance to stretching and compressing), with strong bonds typically being stiffer than weak bonds.
- the bond strength increases as the multiplicity of a bond increases, decreases as the number of lone pairs on the neighboring atoms increases, and decreases as the atomic radii increase. Bonds are strengthened by resonance.
- a bond length is the distance between the centers of two atoms joined by a covalent bond. It corresponds to the internuclear distance at the potential-energy minimum for the two atoms. Bond lengths affect the overall size and shape of a molecule.
- the bond strength decreases down each group as the atoms increase in size
- bonds between heavy atoms tend to be longer than those between light atoms because heavier atoms have larger radii than lighter ones. Multiple bonds are shorter than single bonds between the same two elements, because the additional bonding electrons attract the nuclei more strongly and pull the atoms closer together.
- for bonds between atoms of the same two elements, the stronger the bond the shorter it is. Thus, a carbon triple bond is both stronger and shorter than a carbon double bond.
- each bond makes a characteristic contribution called its covalent radius, to the length of the bond. A bond length is approximately the sum of the covalent radii of the two atoms.
- the covalent radius of an atom taking part in a multiple bond is smaller than that for a single bond of the same atom.
- covalent radii typically decrease from left to right across a period. So, as nuclear charge goes up, the atomic radius decreases. The reason is the same as for atomic radii: the increasing effective nuclear charge draws in the electrons and makes the atom more compact. Like atomic radii, covalent radii increase down a group because, in successive periods, the valence electrons occupy shells that are more distant from the nucleus and are better shielded by the inner core of electrons.
- the covalent radius of an atom is the contribution it makes to the length of the covalent bond; covalent radii are added together to estimate the lengths of bonds in molecules.