Chapter 6

Chapter 6

Chapter 6 Chemical Bonding Introduction Bonding to Section 1 Chemical

Chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together Why are most atoms bonded to each other?

chemically As independent particles, they are at relatively high potential energy Nature, however, favors arrangements in which potential energy is minimized This means that most atoms are less stable existing by themselves than when they are combined By bonding with each other, atoms decrease in potential energy, thereby

creating more stable arrangements of matter Types of Chemical Bonding Bond valence electrons rearranged to make atom more stable Way they are rearranged depends on type of bond Ionic bonding chemical bonding that results from the electrical attraction between large numbers of cations and

anions Covalent bonding results from the sharing of electron pairs between two atoms In purely covalent bond, electrons shared equally between two atoms Ionic or Covalent? Bonding is rarely purely one or the other

Depending on how strongly the atoms attract electrons, falls somewhere between Electronegativity (EN) atoms ability to attract electrons Degree of bonding between atoms of 2 elements being ionic or covalent estimated by calculating difference in elements ENs Example Fluorines EN = 4.0, Cesiums EN

= 0.7 4.0-0.7 = 3.3 According to table, F-Cs is ionic The greater the difference, the more ionic the bond Bonding between atoms with EN difference of less than or equal to () 1.7 has ionic

character less than or equal to () 50% Classified as covalent Bonding between atoms of same element is completely covalent Nonpolar-covalent Bonds

H-H bond has 0% ionic character Nonpolar-covalent bond a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge 0-5% ionic character (0-0.3 EN difference) is nonpolar-covalent bond

Polar-covalent Bonds Bonds that have significantly different Ens, electrons more strongly attracted by moreEN atom These bonds are polar they have an uneven distribution of charge Covalent bonds with 5-50% ionic character (0.3-1.7 EN difference) are polar

+ Polar-covalent bond covalent bond in which the bonded atoms have an unequal attraction for the shared electrons Sample Problem

Use electronegativity differences to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative? Bonding between sulfur

and Hydrogen EN differenc e Bond type Morenegative atom

2.52.1=0.4 Polarcovalent Sulfur Cesium 2.50.7=1.8 Ionic

Sulfur Chlorine 3.02.5=0.5 Polarcovalent chlorine Practice Problem

Use electronegativity differences to classify bonding between chlorine, Cl, and the following elements: calcium, Ca; oxygen, O; and bromine, Br. Indicate the more-negative atom in each pair. Bonding between chlorine

and Calcium EN differenc e Bond type Morenegative atom

3.01.0=2.0 Ionic Chlorine Oxygen 3.53.0=0.5 Polarcovalent

Oxygen Bromine 3.02.8=0.2 Nonpolarcovalent Chlorine Section 2 Covalent

Bonding and Molecular Compounds Many chemical compounds are molecules Molecule neutral group of atoms that are held together by covalent bonds Single molecule of compound is individual unit Capable of existing on its own May consist of 2 or more atoms of same element or two or more different atoms Molecular compound chemical compound whose simplest units are molecules

Formation of Covalent Bond Bonded atoms have lower potential energy than unbonded atoms At large distance atoms dont influence each other Potential energy set at 0 Each H has (+) proton

Nucleus surrounded by (-) electron As atoms near each other, charged particles start to interact Approaching nuclei and electrons are attracted to each other Decrease in total potential energy At the same time, two

nuclei and two electrons repel each other Increase in potential energy The amount of attraction/repulsion depends on how close the atoms are to each other When atoms first see each other, electron-proton attraction stronger than ee or p-p repulsions So atoms drawn to each other and potential energy lowered

Attractive force dominates until a distance is reached where repulsion equals attraction Valley of the curve

Closer the atoms get, potential energy rises sharply Repulsion becomes greater than attraction Characteristics of the Covalent Bond Bonded atoms vibrate a bit

As long as energy stays close to minimum they stay covalently bonded Bond length the distance between two bonded atoms at their minimum potential energy (average distance between two bonded atoms) To form covalent bond, hydrogen atoms

need to release energy Amount of energy equals difference between potential energy at zero level (separated atoms) and at bottom of valley (bonded atoms) Same amount of energy must be added to separate bonded atoms

Bond energy energy required to break a chemical bond and form neutral isolated atoms Units of bond energy usually kJ/mol Indicates energy required to break one mole of bonds in isolated molecules

Ex. 436 kJ/mol is energy needed to break HH bonds in 1 mol hydrogen molecules and form 2 mol of separated H atoms Bond lengths and bond energies vary with the types of atoms that have combined All individual H atoms contain single, unpaired e- in 1s orbital Sharing allows electrons to experience

effect of stable electron configuration of helium, 1s2 The Octet Rule Noble-gas atoms have minimum energy existing on their own b/c of electron configurations Outer orbitals completely full

Other atoms full orbitals by sharing electrons Bond formation follows octet rule chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level

Example: Bonding of Fluorine 2 F atoms bond to form F2 7 e- in highest energy level

Example: HCl Exceptions to Octet Rule Most main-group elements form covalent bonds according to octet rule Ex. H-H only 2 electrons Boron, B, has 3 valence electrons ([He]2s22p1) Boron tends to form bonds where it is surrounded by 6 e- (e- pairs) Others can be surrounded by more than 8

when bonding to highly electronegative elements Electron Dot Notation Objective 01 I will know how to draw Lewis structures for molecules containing single bonds and multiple bonds CLE 3221.3.1 Investigate chemical bonding

To keep track of valence electrons, it is helpful to use electron dot notation electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the elements symbol Inner-shell electrons NOT shown Sample Problem 1

Write the electron-dot notation for hydrogen. A hydrogen atom has only one occupied energy level, the n=1 level, which contains a single electron. So, e-dot notation is written as

H Sample Problem 2 Write the e-dot notation for nitrogen. Group notation for nitrogens family is ns2np3 which means nitrogen has 5 valence electrons. E-dot notation is written as

Drawing Lewis Structures Objective 02 I will know how to draw Lewis structures for molecules containing single bonds and multiple bonds CLE 3221.3.1 Investigate chemical bonding Lewis structures formulas in which atomic symbols represent nuclei and innershell electrons, dot-pairs or dashes between

two atomic symbols represent electron pairs in covalent bonds, and dots next to only one atomic symbol represent unshared electrons Common to write Lewis structures that show only shared e- using dashes Structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule F-F

H - Cl Drawing Lewis Structures Step #1 Determine the number of atoms of each element present in the molecule

Drawing Lewis Structures Step #2 Write the electron dot notation for each type of atom Drawing Lewis Structures Step #3 Count

all valence electrons- these must be shown in the final diagram Drawing Lewis Structures Step #4 Arrange the atoms to form a skeleton structure for the molecule, and connect

the atoms by electron-pair bonds If carbon is present, it is the central atom. Otherwise, the leastelectronegative atom is central (except for hydrogen, which is never central) Drawing Lewis Structures Step #5 Add

unshared pairs of electrons to each nonmetal atom (except hydrogen) such that each is surrounded by eight electrons Drawing Lewis Structures Step #6 Count the electrons in the Lewis structure to be sure that the number of valence electrons used equals the

number available If too many have been used subtract one or more lone pairs until the number of valence electrons is correct. Use the lone pairs to make double or triple bonds Sample Problem Draw the Lewis structure of

iodomethane, CH3I. 1. Determine type and number of atoms in molecule. 1 C, 1 I, 3 H 2. Write the e-dot notation for each type of atom in the molecule. 3. Determine the total number of valence ein the atoms to be combined. C 1 x 4e- = 4eI 1 x 7e- = 7eH 3 x 1e- = 3e14e-

4. Arrange the atoms to form a skeleton structure for the molecule If carbon is present, it is the central atom Otherwise, the least-electronegative atom is central (except for hydrogen which is NEVER central) Then connect the atoms by electron-pair bonds. 5. Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is

surrounded by 8 electrons. 6. Count the electrons in the structure to be sure that the number of valence e- used equals the number available. Be sure the central atom and other atoms besides H have an octect. There are eight e- in the four covalent bonds and six e- in the three unshared pairs, giving the correct total of 14 valence electrons

Practice Problem Draw the Lewis structure of ammonia, NH3. Practice Problem Draw the Lewis structure for hydrogen sulfide, H2S.

Draw Lewis Structures for the following molecules: CF 4 SeH NI 2 3

CH 3 Cl Multiple Covalent Bonds Atoms of same elements (especially C, N and O) can share more than one e- pair Double bond covalent bond made by the sharing of two pairs of e- between two

atoms Shown by two side-by-side pairs of dots or two parallel dashes Triple bond covalent bond made by sharing of 3 pairs of e- between 2 atoms Ex. N 2 Each N has 5 valence Each N shares 3 e- with other

Multiple bonds double and triple bonds Double bonds have higher bond energies and are shorter than single bonds

Triple bonds have higher bond energies and are shorter than double bonds Practice Problem Draw the Lewis structure for methanal, CH2O, which is also known as formaldehyde. Practice Problem

Draw the Lewis structure for carbon dioxide. Practice Problem Draw the Lewis structure for hydrogen cyanide, which contains one hydrogen atom, one carbon atom, and one nitrogen atom.

Resonance Structures Some molecules/ions cannot be represented correctly by single Lewis structure Ex. Ozone (O ) 3 Each structure has one single and one double bond

Chemists used to think ozone spends time alternating or resonating between two structures Now know that actual structure is something like an average between the two Resonance bonding in molecules or ions that cannot be correctly represented by a single Lewis structure

To indicate resonance, double-headed arrow placed between resonance structures Section 3 Ionic Bonding and Ionic Compounds Ionic Bonding Ionic compound composed of positive and negative ions that are combined so that the numbers of positive and negative

charges are equal Most exist as crystalline solids, a 3-D network of (+) and (-) ions mutually attracted to one another Different from molecular compound b/c ionic compound not made of independent, neutral units

Chemical formula represents simplest ratio of compounds combined ions that give electrical neutrality Chemical formula of ionic compound shows ratio of ions present in ANY sample of ANY size

Formula unit simplest collection of atoms from which an ionic compounds formula can be recognized Ex. NaCl is formula unit for sodium chloride One sodium cation and one chlorine anion Ratio of ions in formula depends on charges of ions combined

Ex. Calcium and fluorine Ca2+ F1- = total +1 So need 2 F1- to equal +2+(-2) = 0 Formula unit is CaF2 The Octet Rule

Noble-gas atoms have minimum energy existing on their own b/c of electron configurations Outer orbitals completely full Other atoms full orbitals by gaining, losing or sharing electrons

octet rule chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level Formation of Ionic Compounds E-dot notation can be used to demonstrate changes that take place in ionic bonding Do not usually form by combination of

isolated ions Sodium readily gives up 1 e Chlorine readily accepts 1e Characteristics of Ionic Bonding In ionic crystals, ions minimize potential energy by combining in orderly arrangement called a crystal lattice

Attractive forces: between oppositely charged ions (cations and anions) and between nuclei and electrons Repulsive forces: between like-charged ions and between electrons

Crystal lattice structure represents balance between these two forces Within arrangement, each Na+ is surrounded by 6 Cl At the same time, each Cl- is surrounded by 6 Na+

3-D arrangements of ions and strengths of attraction are different with sizes and charges of ions and number of ions of different charges Ex. CaF2, there are 2 anions for each cation Each Ca2+ is surrounded by 8 F Each F- is surrounded by 4 Ca2+

Lattice Energy To compare bond strengths in ionic compounds, chemists compare amounts of energy released when separated ions in gas form crystalline solid Lattice energy energy released when one mole of an ionic compound is formed

from gaseous ions Comparison of Ionic and Molecular Compounds Force that holds ions together in ionic compounds is very strong overall between opposite charges

Molecular compound bonds making up each molecule also strong, but forces between molecules not strong Because of bond strength difference, molecular compounds melt at lower temperatures

Ionic compounds have higher melting and boiling points Ionic compounds are hard but brittle Slight shift of one row of ions causes large buildup of repulsive forces

Repulsive forces make layers split completely In solid state ions cannot move compounds are not electrical conductors

Molten state ions can move freely and can carry electric current Many ionic compounds dissolve in water Attraction to water molecules overcomes attraction to each other

Polyatomic Ions Certain atoms bond covalently to each other to form group of atoms that has molecular AND ionic characteristics Polyatomic ion a charged group of covalently bonded atoms

Lewis Structures of Polyatomic Ions Polyatomic ions combine with ions of opposite charge to form ionic compounds To find Lewis structure, follow previous instructions except If ion is negative, add to the total number of valence electrons a number of e- same as the ions negative charge If ion positive, subtract same number of eas the positive charge

Section 4 Metallic Bonding Metallic Bonding is Different Metals have unique property of highly movable electrons (why they conduct electricity so well) In molecular compounds e- cannot move, held in shared bond

In ionic compounds, e- cannot move, held to individual ions Metallic-Bond Model Highest energy levels of most metal atoms only occupied by few e-

Ex. s-block metals have one or two valence e- where all 3 p orbitals are empty d-block metals have many empty d orbitals just below highest energy level Overlapping Orbitals Within metal, empty orbitals in outer energy levels overlap Allows outer e- to move freely

e- are delocalized do not belong to any one atom Metallic bonding chemical bonding that results from attraction between metal atoms and surrounding sea of electrons Metallic Properties Freedom of e- to move around causes high electrical and thermal conductivity b/c many orbitals separated by very small energy differences, metals can absorb wide

range of light frequencies Absorption of light excites e- to higher energy levels e- immediately fall back down to lower levels, giving off light (why metals are shiny) Malleability ability of a substance to be hammered or beaten into thin sheets Ductility ability of a substance to be pulled into wires

Both possible because of structure, one line of metal atoms can slide without breaking bonds Not possible with ionic crystal structures Metallic Bond Strength Bond strength varies with nuclear charge of metal atoms and number of e- in metals esea Both factors reflected as heat of vaporization

When metal vaporized, bonded atoms in solid state converted to individual atoms in gas state Higher heat of vaporization, higher bond strength Section 5 Molecular Geometry Molecular Geometry Properties of molecules depend on bonding

of atoms and the 3-Dimensional arrangement of molecules atoms in space Polarity of each bond, along with geometry of molecule, determines molecular polarity uneven distribution of molecular charge Strongly influences forces that act BETWEEN molecules VSEPR Theory

Diatomic molecules must be linear (only two atoms) To predict geometries of more complex molecules, consider locations of all e- pairs surrounding bonded atoms

This is basis of VSEPR Valence-shell, electron-pair repulsion VSEPR theory repulsion between the sets of valence-level e- surrounding an atom causes these sets to be oriented as far apart as possible

How does this account for molecular shape? Lets consider only molecules with no unshared valence e- on central atom Ex. BeF2 Be doesnt follow octect rule Be forms covalent bond with each F atom Surrounded by only two electron pairs it shares with F atoms

According to VSEPR, shared pairs oriented as far away from each other as possible Distance between e- pairs maximized if bonds to F are on opposite sides of Be, 180 apart

So, all 3 atoms lie in straight line molecule is linear If we represent central atom in molecule by A and atoms bonded to A are represented by B then BeF2 is an example of an AB2 molecule

AB2 is linear What would AB3 look like? The 3 A-B bonds stay farthest apart by pointing to corners of equilateral triangle, giving 120 between bonds

= trigonal-planar geometry AB4 molecules following octect rule by sharing 4 e- pairs with B atoms Distance between e- pairs maximized if each A-B bond points to one of 4 corners of tetrahedron (tetrahedral geometry) Angle is 109.5 Sample Problem

Use VSEPR theory to predict the molecular geometry of aluminum chloride, AlCl3. This molecule is an exception to the octet rule because in this case Al forms only three bonds

Aluminum trichloride is an AB3 type of molecule Therefore, according to VSEPR theory, it should have trigonal-planar geometry Practice Problem

Use VSEPR theory to predict the molecular geometry of the following

molecules: a. HI linear b. CBr4 tetrahedral c. AlBr3 Trigonal-planar d. CH2Cl2 tetrahedral VSEPR and Unshared ePairs

Ammonia, NH3, and water, H2O, are examples of molecules where central atom has both shared and unshared e- pairs How does VSEPR account for the geometries? Lewis structure of ammonia shows in addition to 3 e- pairs it shares with 3 H

atoms, the central N has one unshared pair of e VSEPR theory says that lone pair occupies space around N atom just as bonding pairs do So, as an AB molecule, e- pairs maximize 4 separation by assuming 4 corners of tetrahedron Lone pairs occupy space but description of shape of molecule refers to positions of

atoms only So, molecular geometry of ammonia molecule is pyramid with triangular base General formula is AB3E E is unshared e- pair

Water molecule has 2 unshared e- pairs It is AB E molecule 2 2 A (O) is at center of tetrahedron 2 corners occupied by B (H) Other 2 corners occupied by E (unshared e-)

Molecular Shape Atoms bonded to central atom Lone pairs of electrons Bond angle

Linear 2 0 180 Bent or Angular 2

1 Less than 120 Trigonalplanar 3 0

120 Tetrahedral 4 0 109.5 Molecular Shape

Atoms bonded to central atom Lone pairs of electrons Bond angle Trigonalpyramidal

3 1 Less than 109.5 Bent or Angular 2

2 Less than 109.5 Trigonalbipyramidal 5 0 90, 120, and

80 Octahedral 6 0 90 and 180 Sample Problem

Use VSEPR theory to predict the shape of a molecule of carbon dioxide, CO2. Intermolecular Forces Intermolecular forces forces of attraction between molecules Vary in strength Generally weaker than bonds that hold

molecules together (covalent, ionic) Molecular Polarity and DipoleDipole Forces Strongest intermolecular forces exist between polar molecules Polar molecules act as tiny dipoles b/c of

uneven charge distribution Dipole created by equal but opposite charges that are separated by a short distance Direction of dipole is from dipoles positive pole to its negative pole

Represented by arrow with head pointing toward negative pole and crossed tail pointing toward positive pole H Cl Cl more electronegative, and so is negative end

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