Chemical Bonding I: Basic Concepts

Chemical Bonding I: Basic Concepts
Chapter 9

Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that participate in chemical bonding. Group # of valence e- e- configuration 1A 1 ns1 2A 2 ns2 3A 3 ns2np1 4A 4 ns2np2 5A 5 ns2np3 6A 6 ns2np4 7A 7 ns2np5

Lewis Electron-Dot Symbols
A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element. Note that the group number indicates the number of valence electrons. Na . Si : P S Mg Al Cl Ar Group I Group II Group VII Group VIII Group VI Group IV Group V Group III 2

Lewis symbol (Lewis structure) - a way to represent atoms (and their bonds) using the element symbol and valence electrons as dots.

Chemical Bonding Chemical Bond - the force of attraction between any two atoms in a compound. Interactions involving valence electrons are responsible for the chemical bond.

Describing Ionic Bonds
An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions. This type of bond involves the transfer of electrons from one atom (usually a metal) to another (usually a nonmetal). The number of electrons lost or gained by an atom is determined by its need to be “isoelectronic” with a noble gas. 2

Let’s examine the formation of NaCl
IONIC BONDING Let’s examine the formation of NaCl Na + Cl  NaCl Chlorine has a high electron affinity. When chlorine gains an electron, it gains the Ar configuration. Sodium has a low ionization energy it readily loses this electron . When Sodium loses the electron, it gains the Ne configuration. Na  Na+ + e-

Essential Features of Ionic Bonding
Atoms with low I.E. and low E.A. tend to form positive ions. Atoms with high I.E. and high E.A. tend to form negative ions. Ion formation takes place by electron transfer. The ions are held together by the electrostatic force of the opposite charges. Reactions between metals and nonmetals (representative) tend to be ionic.

Such noble gas configurations and the corresponding ions are particularly stable. The atom that loses the electron becomes a cation (positive). The atom that gains the electron becomes an anion (negative). 2

Consider the transfer of valence electrons from a sodium atom to a chlorine atom. e- The resulting ions are electrostatically attracted to one another. The attraction of these oppositely charged ions for one another is the ionic bond. 2

Electron Configurations of Ions
As metals lose electrons to form cations and establish a “noble gas” configuration, the electrons are lost from the valence shell first. For example, magnesium generally loses two electrons from its 3s subshell to look like neon. [Ne]3s2 [Ne] 2

Electron Configurations of Ions
Transition metals also lose electrons from the valence shell first, which is not the last subshell to fill according to the aufbau sequence. For example, zinc generally loses two electrons from its 4s subshell to adopt a “pseudo”-noble gas configuration. [Ar]4s23d10 [Ar]3d10 2

The Ionic Bond - - - Li+ F Li + F 1s22s1 1s22s22p5 1s2 1s22s22p6 [He]
[Ne] Li Li+ + e- e- + F - F - Li+ + Li+

Lewis Electron-Dot Formulas
A Lewis electron-dot formula is an illustration used to represent the transfer of electrons during the formation of an ionic bond. As an example, let’s look at the transfer of electrons from magnesium to fluorine to form magnesium fluoride. 2

Lewis Electron-Dot Formulas
The magnesium has two electrons to give, whereas the fluorines have only one “vacancy” each. : F . Mg [ F ] - 2+ Consequently, magnesium can accommodate two fluorine atoms. 2

Metals and nonmetals usually react to form ionic compounds.
I. Ionic Compounds Metals and nonmetals usually react to form ionic compounds. The metals are the cations and the nonmetals are the anions. The cations and anions arrange themselves in a regular three-dimensional repeating array called a crystal lattice.

Here is the crystal lattice of sodium chloride (table salt.)
Formula - the smallest whole number ratio of ions in the crystal. The formula for sodium chloride is NaC.

Writing Formulas of Ionic Compounds
Determine the charge of the ions (usually can be obtained from the group number.) Cations and anions must combine to give a formula with a net charge of zero It must have the same number of positive charges as negative charges.

Predicting Formulas Predict the formula of the ionic compounds formed from the following combination of ions: 1. sodium and sulfur 2. magnesium and oxygen 3. aluminum and oxygen 4. barium and fluorine

Energy Involved in Ionic Bonding
The transfer of an electron from a sodium atom to a chlorine atom is not in itself energetically favorable; it requires 147 kJ/mol of energy. However, 493 kJ of energy is released when these oppositely charged ions come together. An additional 293 kJ of energy is released when the ion pairs solidify. This “lattice energy” is the negative of the energy released when gaseous ions form an ionic solid. The next slide illustrates this. 2

Energy Involved in Ionic Bonding
2

Electrostatic (Lattice) Energy
Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. Q+ is the charge on the cation E = k Q+Q- r Q- is the charge on the anion r is the distance between the ions cmpd lattice energy MgF2 MgO 2957 3938 Q= +2,-1 Q= +2,-2 Lattice energy (E) increases as Q increases and/or as r decreases. LiF LiCl 1036 853 r F- < r Cl-

Born-Haber Cycle for Determining Lattice Energy
DHoverall = DH1 + DH2 + DH3 + DH4 + DH5 o

Solar Evaporation for Salt
Chemistry In Action: Sodium Chloride Mining Salt Solar Evaporation for Salt

Covalent Bonds When two nonmetals bond, they often share electrons since they have similar attractions for them. This sharing of valence electrons is called the covalent bond. These atoms will share sufficient numbers of electrons in order to achieve a noble gas electron configuration (that is, eight valence electrons). 2

Why should two atoms share electrons?
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? 7e- 7e- 8e- 8e- F F + F Lewis structure of F2 lone pairs F single covalent bond single covalent bond F

not a part of a massive three dimensional crystal structure.
II. Covalent Compounds Covalent compounds are usually formed from nonmetals. Molecules - compounds characterized by covalent bonding. not a part of a massive three dimensional crystal structure.

Let’s look at the formation of H2 H + H  H2
COVALENT BONDING Let’s look at the formation of H2 H + H  H2 Each hydrogen has one electron in it’s valance shell. If it were an ionic bond it would look like this: However, both hydrogen atoms have the same tendency to gain or lose electrons. Both gain and loss will not occur.

Instead, each atom gets a noble gas configuration by sharing electrons.
Each Hydrogen atom now has two electrons around it and has a He configuration The shared electrons pair is a Covalent Bond

H2, N2, O2, F2, Cl2, Br2, I2 Features of Covalent Bonds
Covalent bonds tend to form between atoms with similar tendency to gain or lose electrons. The diatomic elements have totally covalent bonds (totally equal sharing.) H2, N2, O2, F2, Cl2, Br2, I2 Each fluorine has eight electrons around it. Ne’s configuration.

Lewis structure of water
single covalent bonds 2e- 8e- 2e- H + O + H O H or Double bond – two atoms share two pairs of electrons 8e- 8e- 8e- double bonds O C or O C double bonds Triple bond – two atoms share three pairs of electrons triple bond 8e- N 8e- or N triple bond

Lengths of Covalent Bonds
Bond Lengths Triple bond < Double Bond < Single Bond

Polar Covalent Bonds A polar covalent bond is one in which the bonding electrons spend more time near one of the two atoms involved. When the atoms are alike, as in the H-H bond of H2 , the bonding electrons are shared equally (a nonpolar covalent bond). When the two atoms are of different elements, the bonding electrons need not be shared equally, resulting in a “polar” bond. 2

Polar Covalent Bonding and Electronegativity
1 The Polar Covalent Bond Ionic bonding involves the transfer of electrons. Covalent bonding involves the sharing of electrons. Polar covalent bonding - bonds made up of unequally shared electron pairs.

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms H F electron rich region electron poor region e- poor e- rich F H d+ d-

The electrons spend more time with fluorine. This sets up a polar bond
somewhat positively charged somewhat negatively charged These two electrons are not shared equally. The electrons spend more time with fluorine. This sets up a polar bond A truly covalent bond can only occur when both atoms are identical. Electronegativity is used to determine if a bond is polar and who gets the electrons the most.

Polar Covalent Bonds : : : :
For example, the bond between carbon and oxygen in CO2 is considered polar because the shared electrons spend more time orbiting the oxygen atoms. : d+ : d- d- : : The result is a partial negative charge on the oxygens (denoted d-) and a partial positive charge on the carbon (denoted d+) 2

Electron Affinity - measurable, Cl is highest
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electron Affinity - measurable, Cl is highest X (g) + e X-(g) Electronegativity - relative, F is highest

Polar Covalent Bonds Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, electronegativity increases from the lower-left corner to the upper-right corner of the periodic table. The current electronegativity scale, developed by Linus Pauling, assigns a value of 4.0 to fluorine and a value of 0.7 to cesium. 2

The Electronegativities of Common Elements

Variation of Electronegativity with Atomic Number

Polar Covalent Bonds The absolute value of the difference in electronegativity of two bonded atoms gives a rough measure of the polarity of the bond. When this difference is small (less than 0.5), the bond is nonpolar. When this difference is large (greater than 0.5), the bond is considered polar. If the difference exceeds approximately 1.8, sharing of electrons is no longer possible and the bond becomes ionic. 2

Classification of bonds by difference in electronegativity
Bond Type Covalent  2 Ionic 0 < and <2 Polar Covalent Increasing difference in electronegativity Covalent share e- Polar Covalent partial transfer of e- Ionic transfer e-

Which would be more polar, a H-F bond or a H-Cl bond?
The greater the difference in electronegativity between two atoms, the greater the polarity of a bond. Which would be more polar, a H-F bond or a H-Cl bond? H-F … = 1.9 H-Cl… = 0.9 Therefore, the HF bond is more polar than the HCl bond.

Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent

Drawing Lewis Structures on Molecules and Polyatomic Ions
Lewis Structure Guidelines 1. Use chemical symbols for the various elements to write the skeletal structure of the compound. the least electronegative atom will be placed in the central position, hydrogen and fluorine occupy terminal positions, carbon often forms chains of carbon-carbon covalent bonds. 5

2. Determine the total number of valence electrons associated with each atom in the compound.
for polyatomic cations, subtract one electron for every positive charge; for polyatomic anions, add one electron for every negative charge. 3. Connect the central atom to each of the surrounding atoms using electron pairs. Then give each atom an octet. Remember, hydrogen needs only two electrons

Count the number of electrons you have and compare to the number you used.
If they are the same, you are finished. If you used more electrons than you have add a bond for every two too many you used. Then give every atom an octet. If you used less electrons than you have….(see later when discuss exceptions to the octet rule) 5. Check that all atoms have the octet rule satisfied and that the total number of valance electrons are used.

Covalent Bonds The tendency of atoms in a molecule to have eight electrons in their outer shell (two for hydrogen) is called the octet rule. 2

Lewis Structures You can represent the formation of the covalent bond in H2 as follows: H . : + This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots. 2

Lewis Structures The shared electrons in H2 spend part of the time in the region around each atom. : H In this sense, each atom in H2 has a helium configuration. 2

. : : + Lewis Structures H Cl H Cl
The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way. H . : Cl + : H Cl Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar). 2

Lewis Structures Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures. An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared). bonding pair lone pair : H Cl 2

Coordinate Covalent Bonds
When bonds form between atoms that both donate an electron, you have: A . : B + It is, however, possible that both electrons are donated by one of the atoms. This is called a coordinate covalent bond. A : B + 2

Multiple Bonds In the molecules described so far, each of the bonds has been a single bond, that is, a covalent bond in which a single pair of electrons is shared. It is possible to share more than one pair. A double bond involves the sharing of two pairs between atoms. or C : H 2

Multiple Bonds Triple bonds are covalent bonds in which three pairs of electrons are shared between atoms. C or H : ::: 2

Writing Lewis Dot Formulas
The following rules allow you to write electron-dot formulas even when the central atom does not follow the octet rule. To illustrate, we will draw the structure of PCl3, phosphorus trichloride. 2

Step 1: Total all valence electrons in the molecular formula. That is, total the group numbers of all the atoms in the formula. 5 e- (7 e-) x 3 26 e- total For a polyatomic anion, add the number of negative charges to this total. For a polyatomic cation, subtract the number of positive charges from this total. 2

Step 2: Arrange the atoms radially, with the least electronegative atom in the center. Place one pair of electrons between the central atom and each peripheral atom. Cl Cl P Cl 2

Step 3: Distribute the remaining electrons to the peripheral atoms to satisfy the octet rule. : Cl Cl P Cl 2

Step 4: Distribute any remaining electrons to the central atom. If there are fewer than eight electrons on the central atom, a multiple bond may be necessary. : : : : : : Cl Cl : P : : Cl : 2

Try drawing Lewis dot formulas for the following covalent compound. SCl2 20 e- total Cl S 16 e- left : 4 e- left : 2

Try drawing Lewis dot formulas for the following covalent compound. COCl2 24 e- total 18 e- left : 0 e- left Cl C O 2

Note that the carbon has only 6 electrons. One of the oxygens must share a lone pair. COCl2 24 e- total 18 e- left : 0 e- left Cl C O 2

Note that the carbon has only 6 electrons. One of the oxygens must share a lone pair. COCl2 24 e- total 18 e- left : O : 0 e- left C Note that the octet rule is now obeyed. : : Cl Cl 2

5 + (3 x 7) = 26 valence electrons
Write the Lewis structure of nitrogen trifluoride (NF3). Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons F N

4 + (3 x 6) + 2 = 24 valence electrons
Write the Lewis structure of the carbonate ion (CO32-). Step 1 – C is less electronegative than O, put C in center Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e- 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24 O C

Let's Practive Lewis Structures
Using the guidelines presented, write Lewis structures for the following: 1. H2O 2. NH3 3. CO2 4. NH4+ 5. CO32- 6. N2

triple bond > double bond > single bond
Bond energy - the amount of energy required to break a bond holding two atoms together. triple bond > double bond > single bond Bond length - the distance separating the nuclei of two adjacent atoms. single bond > double bond > triple bond

( ) - - Two possible skeletal structures of formaldehyde (CH2O) H C O
An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. formal charge on an atom in a Lewis structure = total number of valence electrons in the free atom - total number of nonbonding electrons - 1 2 total number of bonding electrons ( ) The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

( ) - -1 +1 H C O = 1 2 = 4 - 2 - ½ x 6 = -1 = 6 - 2 - ½ x 6 = +1
C – 4 e- O – 6 e- 2H – 2x1 e- 12 e- 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 H C O formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons ( ) total number of valence electrons in the free atom - total number of nonbonding electrons formal charge on C = 4 - 2 - ½ x 6 = -1 formal charge on O = 6 - 2 - ½ x 6 = +1

( ) - H C O = 1 2 = 4 - 0 - ½ x 8 = 0 = 6 - 4 - ½ x 4 = 0 C – 4 e-
H C O C – 4 e- O – 6 e- 2H – 2x1 e- 12 e- 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons ( ) total number of valence electrons in the free atom - total number of nonbonding electrons formal charge on C = 4 - 0 - ½ x 8 = 0 formal charge on O = 6 - 4 - ½ x 4 = 0

Formal Charge and Lewis Structures
For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. Lewis structures with large formal charges are less plausible than those with small formal charges. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH2O? H C O -1 +1 H C O

In certain instances, more than one feasible Lewis structure can be illustrated for a molecule. For example, : : H C N H N C or The concept of “formal charge” can help discern which structure is the most likely. 2

The formal charge of an atom is determined by subtracting the number of electrons in its “domain” from its group number. 1 e- 4 e- 5 e- “domain” electrons group number I IV V : : H C N H N C or The number of electrons in an atom’s “domain” is determined by counting one electron for each bond and two electrons for each lone pair. 2

The most likely structure is the one with the least number of atoms carrying formal charge. If they have the same number of atoms carrying formal charge, choose the structure with the negative formal charge on the more electronegative atom. or H C N : formal charge +1 -1 In this case, the structure on the left is most likely correct. 2

Lewis Structures and Resonance
Write the Lewis structure at CO32- on your paper. If you look at the people around you they probably put the double bond in different places. Who is right? You all are.

A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. O + - O + - What are the resonance structures of the carbonate (CO32-) ion? O C - O C - O C -

Delocalized Bonding: Resonance
According to theory, one pair of bonding electrons is spread over the region of all three atoms. O This is called delocalized bonding, in which a bonding pair of electrons is spread over a number of atoms. 2

Exceptions to the Octet Rule
Although many molecules obey the octet rule, there are exceptions where the central atom has more than eight electrons. Generally, if a nonmetal is in the third period or greater it can accommodate as many as twelve electrons, if it is the central atom. These elements have unfilled “d” subshells that can be used for bonding. 2

Lewis Structures and Exceptions to the Octet Rule
1. Incomplete Octet - less then eight electrons around an atom other than H. Let’s look at BF3 2. Odd Electron - if there is an odd number of valence electrons it isn’t possible to give every atom eight electrons. Let’s look at NO 3. Expanded Octet - elements in 3rd period and below may have 10 and 12 electrons around it.

The Incomplete Octet Be – 2e- 2H – 2x1e- 4e- BeH2 H Be B – 3e- 3F – 3x7e- 24e- 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24 F B BF3

Odd-Electron Molecules N – 5e- O – 6e- 11e- NO N O The Expanded Octet (central atom with principal quantum number n > 2) S F S – 6e- 6F – 42e- 48e- 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48 SF6

For example, the bonding in phosphorus pentafluoride, PF5, shows ten electrons surrounding the phosphorus. : F : : F : : : F : P F : : : : F : 2

In xenon tetrafluoride, XeF4, the xenon atom must accommodate two extra lone pairs. : F : F : : : Xe : F : F : : : 2

Write the Lewis structures of SH6 and SF4
Expanded octet is the most common exception. Write the Lewis structures of SH6 and SF4

Lewis Lab – class activity

Single bond < Double bond < Triple bond
The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy H2 (g) H (g) + DH0 = kJ Cl2 (g) Cl (g) + DH0 = kJ HCl (g) H (g) + Cl (g) DH0 = kJ O2 (g) O (g) + DH0 = kJ O N2 (g) N (g) + DH0 = kJ N Bond Energies Single bond < Double bond < Triple bond

Average bond energy in polyatomic molecules
H2O (g) H (g) + OH (g) DH0 = 502 kJ OH (g) H (g) + O (g) DH0 = 427 kJ Average OH bond energy = 2 = 464 kJ

Bond Energies (BE) and Enthalpy changes in reactions
Imagine reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products. DH0 = total energy input – total energy released = SBE(reactants) – SBE(products)

H2 (g) + Cl2 (g) HCl (g) 2H2 (g) + O2 (g) H2O (g)

Use bond energies to calculate the enthalpy change for:
H2 (g) + F2 (g) HF (g) DH0 = SBE(reactants) – SBE(products) Type of bonds broken Number of bonds broken Bond energy (kJ/mol) Energy change (kJ) H 1 436.4 F 1 156.9 Type of bonds formed Number of bonds formed Bond energy (kJ/mol) Energy change (kJ) H F 2 568.2 1136.4 DH0 = – 2 x = kJ

Bond Energy To illustrate, let’s estimate the DH for the following reaction. In this reaction, one C-H bond and one Cl-Cl bond must be broken. In turn, one C-Cl bond and one H-Cl bond are formed. 2

Bond Energy simple arithmetic yields DH. 2

WORKED EXAMPLES

Worked Example 9.2

Worked Example 9.4

Worked Example 9.5

Worked Example 9.6

Worked Example 9.7

Worked Example 9.11

Chemical Bonding I: Basic Concepts

Similar presentations

Presentation on theme: "Chemical Bonding I: Basic Concepts"— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

Chemical Bonding I: Basic Concepts

Similar presentations

Presentation on theme: "Chemical Bonding I: Basic Concepts"— Presentation transcript:

Similar presentations

About project

Feedback