Chemical Bonding - Ms. Tabors Classroommstaborsclassroom.weebly.com/.../5._chemical_bonding.docx ·...

Name: ______________________Chemistry 1 Notes

Topic 5 – Chemical BondingTopics:1.2.1 – Compare (qualitatively) the relative strengths of ionic, covalent, and metallic bonds.1.2.2 – Infer the type of bond and chemical formula formed between atoms.1.2.3 – Compare inter- and intra-particle forces.1.2.5 – Compare the properties of ionic, covalent, metallic and network compounds.

TYPES OF BONDS

Atoms form bonds with one another because in doing so, they become more stable! Being stable usually means having a full octet of eight valence electrons. Atoms achieve this full octet by either sharing electrons or gaining and losing electrons – whether

electrons are shared or transferred determines the type of bond formed.

There are 3 main types of bonds between atoms:1) Covalent – two atoms share electrons2) Ionic – attraction between positive and negative ions

(one atom completely loses one or more electrons and another atom gains them)3) Metallic – all metal ions are joined & electrons are free to move throughout the metal

I. Ionic Bonds An ionic bond results from the electrostatic attraction between positive and negative ions.

Ions are charged particles formed when atoms gain or lose electrons. cations = positively (+) charged ions (formed when atoms lose e–) anions = negatively (–) charged ions (formed when atoms gain e–)

The cations and anions in an ionic bond will combine in such a way that their charges are equal and opposite, and so cancel each other out. Thus, even though an ionic compound is composed of these charged ions, the compound remains neutral. The atom which becomes a cation simply gives up its valence e– to the atom which becomes an anion. (If you think about it, all the electrons are still there. They never left the ionic compound, they were simply transferred from the cation to the anion.)

Ionic compounds such as NaCl do not consist of just one Na+ ion bonded to one Cl– ion. Rather, the formula for an ionic compound represents the relative ratios of these ions in a huge lattice of thousands or millions of ions (e.g. – in a crystal of table salt, NaCl, if there are 5,000 Na+ ions, then there are 5,000 Cl– ions). Since the formula NaCl does not represent the actual number of ions in a crystal of table salt, but rather their ratios, it is called a formula unit. The formulas of all ionic compounds are likewise formula units.

As is illustrated here, ionic compounds often exist as crystalline structures, arranged so that the oppositely charge ions are next to each other.

The distance between ions in a crystal represents a balance between the attraction and repulsion of ions (the lowest possible energy = most stable distance).

1

NaCl crystal lattice

♦ A quick video segment showing formation of NaCl:021_SODIUMCHLOR.MOV

http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/021_SODIUMCHLOR.MOV


As we mentioned, elements usually form ions so as to achieve a full octet of eight valence electrons (thus they have a noble gas electron configuration). Here is a look at how atoms or groups of atoms form ions: Group 1 metals: +1 ions (ex: Na+) by losing one valence electron → full octet. Group 2 metals: +2 ions (ex: Ca2+) by losing two valence electrons → full octet. Group 3 (or 13) metals: +3 ions (ex: Al3+) by losing three valence e– → full octet. Group 5 (or 15) non-metals: –3 ions (ex: P3–) by gaining three valence e– → full octet. Group 6 (or 16) non-metals: –2 ions (ex: O2–) by gaining two valence e– → full octet. Group 7 (or 17) non-metals: –1 ions (ex: F–) by gaining one valence e– → full octet. The transition metals can often form more than one ion (multiple oxidation states):

iron forms Fe+2 and Fe+3 ions, copper forms Cu+ and Cu2+ ions, etc.

Clearly, then, the most important stable electron arrangement to remember is that of the noble gases – a full octet. But there are also other electron arrangements that confer some degree of stability to atoms are ions. Consider the valence electron configurations of the following d-block metals:

Take chromium as an example. Its valence electron configuration is 4s13d5. You might expect Cr to be 4s23d4, but it isn’t. Scientists still don’t completely agree on why this happens. But it happens to the 4th and 9th elements in the 3d block. A half-filled d-orbital (5 d electrons) is somewhat more stable than a full s subshell in the

next shell, which helps explain the configurations of Cr and Mo. It takes less energy to maintain an electron in a half-filled d subshell than in a filled s subshell.

Likewise, a full d subshell is also a stable electron arrangement (it takes less energy to maintain an electron in a full d subshell than in a full s subshell). This accounts for the unusual configurations of copper, silver and gold.

This helps to explain the +1 charge possible for some d-block metals. The outer electron configurations of such elements are __s1__d5, and they lose that one s electron to form a +1 ion.

Obviously multiple charges are possible for many of these d-block elements, and the full story here is actually quite complicated. You only need to remember the information we’ve presented here…

Ionic bonds usually occur between a metal and a non-metal, or between a metal and a polyatomic ion.

2


Note the location of the “staircase” on the periodic table to help you remember where the metals, metalloids, and non-metals are.

The metal (having a lower ionization energy) loses its valence electrons, and the non-metal (having a higher electronegativity) gains these electrons.

In the process, both atoms achieve a noble gas electron configuration, and are held together by their opposite charges.

As we’ll soon see in more detail, the type of bond that will be formed between two elements is actually a function of their electronegativities. The greater the difference in electronegativity between two bonded atoms, the more ionic the bond. Consider NaCl and HCl as examples: the electronegativity of chlorine is 3.0, hydrogen is 2.1,

and sodium is 0.9. NaCl has a greater electronegativity difference between the atoms (a difference of 2.1 vs. 0.9 for

HCl), and therefore has the greater ionic character (in fact, HCl is a polar covalent bond).

II. Covalent Bonds A covalent bond forms when atoms share electrons. A covalent bond could be described as the

electrostatic attraction between a pair of electrons and positively charged nuclei.

3


Consider how a covalent bond forms between two hydrogen atoms:

♦ Here is a good animation of this process of bond formation in hydrogen:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/038_H2BondForm.MOV

The term molecule refers to two or more atoms held together by covalent bonds. The term molecule can ONLY be applied to covalent compounds!

(Remember that the equivalent term for an ionic compound is a “formula unit”.) A molecule is a discrete unit. A molecule of water (H2O) contains two hydrogen atoms and one

oxygen atom, and is independent of any water molecules beside it. A molecule containing two atoms is called a diatomic molecule (ex: HCl or H2).

♦ In covalent molecules, the electrons are free to roam in the orbitals of either atom, but spend most of their time between the atoms. This is what holds the bond together and makes the molecule stable.

Similar to ionic bonds, atoms share electrons so that they can both achieve stable noble gas electron configurations. Consider the covalent bonding of hydrogen and oxygen in water, for example:

H has 1 valence electron, while O has 6 valence electrons

Consider carbon dioxide (CO2) as a second example:C has 4 valence e–, while O has 6 valence e–, necessitating double bonding:

4

http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/038_H2BondForm.MOV


As the previous example of CO2 illustrates, it is sometimes necessary to form multiple covalent bonds (double or triple bonds) to satisfy the octets of bonded atoms.

A double bond is a covalent bond where atoms share two pairs of e–

A triple bond is a covalent bond where atoms share three pairs of e–

(only C & N can form triple bonds)

Many carbon-containing compounds (organic compounds, such as ethene & ethyne, above) contain multiple bonds.

LEWIS STRUCTURES & STRUCTURAL FORMULAE

octet rule – Most chemical compounds form so that each atom in the bond can have eight electrons in its highest energy level, which is the most stable electron arrangement (that of an inert or noble gas).

This is illustrated in the way we draw molecules and ions. There are both shared and unshared pairs of electrons which play a role in bonding and the shape

of molecules.

unshared (non-bonding) pair – a pair of electrons not involved in bonding, but belonging to exclusively one atom

ex:

Lewis structures use dots or dashes or Xs to represent shared and unshared electrons.

5


ex:

Lewis structures must show unshared e– pairs! Cl – Cl is not a correct Lewis structure.

There are seven elements which exist in nature as covalently bonded diatomic molecules: H2, N2, O2, F2, Cl2 , Br2, I2

These seven diatomic molecules are often called the “BrINClHOF” seven…

Several of these exist as single bonds (H2, and the halogens), but N2 and O2 exist as triple and double bonds, respectively.

We mentioned previously that some elements do not always adhere to the octet rule. The natural state of some elements involves more than just one or two of the elements joined together in a covalent compound. Consider sulfur (S8), phosphorus (P4), and buckyballs or fullerenes (C60, a form of carbon) shown below:

There is a definite relationship between the # of bonds, bond length, and bond strength: In general, the greater the # of electron pairs (bonds) holding two atoms together, the shorter the

bond length, and the greater the bond strength.

6

Buckyball, with an atom inside of it


Consider the single, double, & triple bonds between carbon atoms as the 1st example from the data table below: multiple bonds are shorter and stronger. (By stronger, we mean they have a higher bond enthalpy – it takes more energy to break the bond and pull the atoms apart.)

Consider also both carbon-oxygen bonds in a carboxylic acid (such as ethanoic acid [a.k.a. acetic acid, or vinegar] ) as shown below.

bond bond length bond enthalpy(nm) (H / kJ mol-1)

C – C 0.154 348C C 0.134 612

C C 0.120 837C – O 0.143 360C O 0.122 743

Refer to your data booklet for additional examples…

However, multiple bonds (while being stronger) also tend to be more reactive – this is very important, so keep this in mind!

Covalent bonds tend to form between non-metals. Remember the “stairstep” on the periodic table, and that hydrogen is a non-metal. However, the true criterion for determining the nature of a bond (as mentioned earlier) is the

difference in electronegativity between bonded atoms:

difference inelectronegativity 4.0 1.7 0.3 0.0

Ionic Polar Non-Polar

If the difference in electronegativity between two bonded atoms is:

between 0.0 – 0.3 non-polar covalent bond – bonding e– are shared equally(This includes the “BrINClHOF” 7 (H2, N2, O2, F2, Cl2 , Br2, I2) which obviously have no e-neg difference between them. Also assume C–C and C–H bonds are effectively non-polar.)

between 0.3 – 1.7 polar covalent bond – the more electronegative atom attracts the shared electron pair more strongly than the less electronegative atom (ex: HCl)

between 1.7 – 4.0 ionic bond – one atom completely loses its valence electrons and the other (more electronegative) atom gains them (ex: NaCl)

7


Keep in mind that bonds are rarely completely covalent or ionic; usually, it’s somewhere in between, depending on the electronegativity difference. We typically talk in terms of the ionic or covalent character of the bond between two atoms. The closer together two elements are on the periodic table, the smaller the difference in

electronegativity between them. This would mean that the bond between them has more covalent character. Two elements farther apart would likely have a greater difference in electronegativity and

thus greater ionic character.

For example, consider these elements (electronegativities in parentheses):H (2.1) Na (0.9) Cl (3.0)

NaCl, with an e-neg difference of 2.1 would have much greater ionic character than HCl, with e-neg difference of 0.9.

In a covalent bond, electron distribution may not be symmetrical around the nuclei. Electrons spend more time next to the atom that attracts them most strongly (the most electronegative atom).

Non-Polar Covalent Bond

PROPERTIES OF IONIC AND COVALENT COMPOUNDS

8

symmetrical electron distribution; equal sharing of electron pair

asymmetrical electron distribution; unequal sharing of electron pair

+ -

Polar Covalent Bond


Covalent bonds in a molecular compound are strong, but ionic bonds in an ionic compound are even stronger. Consider how these physical properties of covalent and ionic compounds reflect the different nature of their bonds:

The stronger forces of attraction in ionic bonds take more energy to overcome, thus giving ionic compounds higher melting and boiling points than covalent compounds.

Similarly, covalent compounds are more likely to be volatile than are ionic compounds. Volatility is the tendency for a liquid to evaporate at room temperature. The weaker forces of attraction between covalent compounds mean the energy available at room temperature is more likely to be sufficient for some of the covalent molecules to evaporate. Consider water, rubbing alcohol, perfumes, gasoline etc. as evidence of this.

Ionic compounds also tend to be much harder and more brittle than covalent compounds, which also has to do with the stronger electrostatic attraction in ionic bonds. This is the reason that ionic compounds are so hard – they don't want to move around much, so they don't bend at all. This also explains the brittleness of ionic compounds. Because of their inflexibility, to break an ionic crystal, we have to give it such a strong whack with a hammer, that the crystal doesn't break in just one spot, but shatters in many places.

Neither ionic nor covalent compounds conduct electricity (moving electrons) as solids. But when you dissolve ionic compounds in water, now the ions are free to move around, and can ferry the electrons with them, facilitating the movement of electrons. The same idea holds true when you melt ionic compounds. As solids, the ions are locked in

place, but when molten, these ions are free to move around. Just like when they are dissolved in aqueous solution, this facilitates the movement of electrons.

Remember that covalent compounds do not contain ions. And when you melt covalent compounds or dissolve them in aqueous solution they have no tendency to form ions. Thus solutions of covalent compounds are not electrolytic (they do not conduct electricity).

9


III. Metallic Bonds Metals readily give up electrons (having low ionization energies), but don’t attract electrons as

strongly as non-metals (which have high electronegativities). In a pure chunk of metal, the valence electrons are delocalized – they do not belong to any one

single atom, but are free to move around metal cations and free electrons. All these mobile, delocalized valence electrons are shared equally by all the atoms, in a sort of

“electron sea”. Thus, a metallic bond results from the attraction between positive metal ions and the surrounding sea of mobile electrons.

The higher the + charge of the metal cation, the stronger the metallic bond formed. (ex: Al atoms forms stronger metallic bonds than Na atoms, b/c Al forms +3 ions, while Na forms +1 ions. The greater strength of the Al metal bond is due in part to a greater # of delocalized valence e-.)

Consider the properties of metals, and how their “electron sea” helps explain some of them: high electrical conductivity – if the delocalized electrons are free to move around, that allows a

current of electricity to flow easily thermal conductivity – metals conduct heat as well

When a metal bar is heated at one end, the electrons at that end increase in kinetic energy and rapidly flow through the metal, transporting heat energy with them.

(By contrast, molecular or ionic compounds can only pass along heat slowly, from one neighboring atom to the next.)

malleability – metals can be shaped or hammered into a flat sheet Again, because of the delocalized electrons, metallic bonding is not directional, but uniform

throughout the solid. One plane of metal ions can slide past one another without encountering any resistance or breaking any bonds.

ductility – metals can be drawn out into a fine wire This property is explained by the same uniform bonding of metal ions mentioned under

malleability. high melting & boiling point – The metallic bond is an extremely strong bond – stronger than

most covalent and ionic bonds. Thus, much energy is required to overcome the metallic bonding, resulting in high melting and boiling points for most metals.

luster – Luster refers to the “shininess” of metals. The delocalized electrons in metals do really absorb or capture light. To the extent that they do absorb light, the electrons quickly return to their original energy level by emitting the energy they absorbed with little dissipation. So, when light strikes a metal surface all visible wavelengths are reflected with little or no alteration in color.

10


IV. VSEPR Molecular Geometry VSEPR – Valence Shell, Electron-Pair Repulsion

VSEPR theory states that the repulsion between valence electron pairs (being negative, they obviously repel each other) causes bonds to be oriented as far apart as possible. Molecules will take the widest bond angle possible between atoms.

When determining VSEPR molecular geometry: How do you determine which atom is in the middle? If the molecule has only single bonds, it

will be whichever atom has most available sites for bonding. The atom written first in the formula of a covalent compound usually is in the middle, unless it’s hydrogen.

If there are more than four valence e–, then there must be unshared electron pairs bent or triangular pyramidal geometry. Use common sense when deciding on the geometry of molecules with multiple bonds.

Draw the following molecules, giving their VSEPR molecular geometries and bond angles:PI3, CO2, SO2, C2H2, C2H4, CO3

2–, NO2–

11

Bond Angle

180°

120°

~118°

109.5°

107°

105°

Description

diatomic, or central atom bonded to 2 peripheral atoms

& central atom has zerounshared e- pairs

central atom bonded to3 peripheral atoms



& central atom has oneunshared e- pair




& central atom has oneunshared e- pair


& central atom has twounshared e- pairs


V. Molecular Polarity Dipoles exist between individual atoms if there is a difference in electronegativity between the two

atoms.

In a covalent bond, electron distribution may not be symmetrical around the nuclei. Electrons spend more time next to the atom that attracts them most strongly (the most electronegative atom).

+ – Non-Polar Covalent Bond Polar Covalent Bond

How do I know when a molecule is polar or non-polar?

non-polar molecules central atom has NO unshared e– pairs central atom is surrounded by identical

atoms, which all have the same electronegativity (ex: CCl4)

polar molecules central atom has an unshared e– pair central atom has at least 2 different

elements surrounding it (atoms which have different electronegativities) (ex: CCl3F)

H2O and NH3 are very polar; but the shape (geometry) of CCl4 and CO2 means that the individual dipoles oppose one another and cancel each other out no molecular dipole.

Consider the molecules you drew on the previous page, & state whether or not they are polar.

12

symmetrical electron distribution; equal sharing of electron pair

asymmetrical electron distribution; unequal sharing of electron pair


VI. Intermolecular Forces of Attraction The forces of attraction between molecules are known as intermolecular forces. These are attractions, not true bonds. They are weaker than covalent, ionic & metallic bonds. Also, remember that these are intermolecular forces. They only occur between molecules (covalent

compounds). The covalent bonds themselves are intramolecular forces.

DIPOLE-DIPOLE FORCES

The strongest intermolecular forces are dipole-dipole forces – forces of attraction between polar molecules. Equal but opposite charges separated by a short distance create a dipole. The dipole arrow points toward the negative pole (more electronegative atom), and the

+ tail points to positive pole (the less electronegative atom).

If you have a solution of polar molecules, the molecules will arrange themselves such that the positive and negative poles are attracted to each other. ex: look at a solution of iodine monochloride…

Dipole-dipole forces will increase the melting & boiling points of a substance, because they cause the polar molecules to “stick together” (like opposite poles of a magnet). Thus, additional energy is required, because these attractions have to be broken before the substance can go from a solid to a liquid (melt) or go from a liquid to a gas (boil).

13


HYDROGEN BONDING

Hydrogen bonding is a particularly strong type of dipole-dipole attraction. It is the strongest of all the intermolecular forces.

Hydrogen bonding is the attraction between a hydrogen atom bonded to a strongly e-neg atom (N, O, F) and the unshared pair of electrons on another strongly e-neg atom (N, O, F). The attraction is usually represented by dashed line (--). Consider water as an example:

Being a strong type of dipole-dipole attraction, hydrogen bonding causes the melting and boiling points of hydrogen-containing compounds to be unusually high – the molecules hold on tightly to each other and don’t want to let go This is a function of the large difference in electronegativity between H and N, O, or F, which

makes the bonds between them highly polar. Compare the melting and boiling points of molecules which exhibit hydrogen-bonding and those

which do not as examples of the difference it makes.

Consider how hydrogen bonding (and other intermolecular forces) affects the boiling points of the following pairs of compounds:

state at room T boiling point (C)H2O liquid 100H2S gas –61CH3–O–CH3 gas –24CH3CH2OH liquid 78NH3 gas –28PH3 gas –90HF gas 21HCl gas –79

VAN DER WAALS’ FORCES (LONDON DISPERSION FORCES )

14


Electrons are constantly in motion in an atom; so, at some given instant, all the electrons could be on one side of the atom, creating a momentary (instantaneous) dipole:

van der Waal’s forces in neighboring helium atoms

- + - + - +

instantaneous dipole instantaneous dipole induced dipole

“before” “after”

van der Waal’s forces are intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles and induced dipoles. Essentially, London dispersion forces occur when the electrons of molecule are attracted to the nucleus of another, neighboring molecule.

The more electrons the stronger the L.D.F. (i.e. – L.D.F. increase as the mass of the atom or molecule increases).

These forces are very weak, because the dipoles created here are temporary (not permanent, based on electronegativity difference, as in dipole-dipole and hydrogen bonding).

L.D.F. act between all molecules, but they are the only intermolecular forces acting on noble gases and non-polar molecular compounds.

The following diagram illustrates the effect of van der Waal’s forces in an H2 molecule.

The polarity of a compound can significantly affect the physical properties of that compound. As an example, consider how a molecule’s solubility is affected by its polarity. Recall that “like dissolves like”.

15


In other words, solvents tend to dissolve solutes with similar polarities. Polar solvents tend to dissolve polar solutes. ex #1: NH3 dissolves in water to make household cleaners – ammonia and water are

both polar, having unshared e– on the central atom of the molecule). ex #2: Water dissolves NaCl, because the ions are attracted to the very polar water

molecules, which play “tug-of-war” with the ionic bonds holding the Na+ and Cl– ions together. NaCl dissolves in water because the pull of hundreds of H2O molecules is stronger than the ionic forces holding Na+ and Cl– ions together. These are called ion-dipole attractions.

Likewise, non-polar solvents tend to dissolve non-polar solutes. ex: I2(s) dissolves in CCl4(l)

16


VII. Macromolecules & Network Covalent Solids ♦ Both macromolecules and network covalent solids are very large covalently-bonded molecules, but

the two terms are not generally interchangeable. These are a class of molecules with unique properties (some similarities to both covalent and ionic compounds).

MACROMOLECULES

♦ The term macromolecule typically refers to synthetic polymers and biopolymers (such as DNA, proteins & hair), as well as large non-polymer molecules such as lipids. Polymers are essentially composed of a relatively small unit (a monomer) that repeats over and

over again.

If we represent a monomer as:

a polymer could be shown as:

or its generalized repeating unit as:

where n is a large number, the value of which can vary considerably.

Examples of synthetic polymers include polyvinyl chloride (PVC) and nylon:

polyvinyl chloride (PVC)

monomer polymer repeating unit

Typical uses: plumbing pipes, vinyl fences and siding, linoleum, raincoats, shower curtains

Poly(chloroethene) is very useful because it is both water repellent and flame retardant (it releases chlorine atoms when burned, which inhibit combustion).

nylon – a polyamide made from two different molecules by condensation reaction

17

Nylon can be used in the manufacture of clothes, fibers, parachutes, toothbrush bristles, balloons, etc.


Examples of biopolymers include DNA, proteins & hair:

18

DNA is a polymer of nucleotides (a ribose sugar, a phosphate group and a nitrogenous base).

Proteins are polymers of amino acids. Examples of proteins include hemoglobin and hair.

Hair is composed primarily of proteins (~90%). These proteins are of a hard, fibrous type known as keratin. Keratin is also the primary component of fingernails, rhinoceros horns, feathers, etc.


NETWORK COVALENT SOLIDS

♦ Network covalent solids are similar to macromolecules but they contain no discrete molecular units (like a monomer). Rather they form a continuous network of covalently bonded atoms. Network covalent solids include allotropes of carbon (graphite, diamonds, fullerenes).

In diamonds, all of the carbon atoms form four identical bonds to neighboring carbons. All of the bonds in this macromolecule are equally strong in each direction (a tetrahedral lattice arrangement). There is no plane of weakness in the molecule, helping to make diamonds the hardest known natural substance.

In graphite, the carbon atoms have three strong bonds to other carbons, all in the same plane (each carbon would have a trigonal planar geometry viewed from above). Each carbon atom has one leftover electron per carbon atom which is delocalized to form very weak bonds between these layers. Since the bonds are so weak, the layers slide fairly freely over each other, making graphite an excellent lubricant.

19


The third allotrope of carbon, C60 (buckminsterfullerenes, or buckyballs, or fullerenes), consists of 60 carbon atoms arranged in hexagons and pentagons to give a geodesic spherical structure similar to a soccer ball.

♦ Silicon and silicon dioxide also exhibit covalent network bonding. Note that silicon and carbon are both in group 4 (group 14).

♦ Also, notice how the hydrogen bonding in ice causes the solid ice to form a structure that resembles the bonding in network solids:

20


21

Chemical Bonding - Ms. Tabors Classroommstaborsclassroom.weebly.com/.../5._chemical_bonding.docx ·...

Documents

Transcript of Chemical Bonding - Ms. Tabors Classroommstaborsclassroom.weebly.com/.../5._chemical_bonding.docx ·...