Chapter 12 Chemical Bonding General info on Chemical

Chapter 12: Chemical Bonding • General info on Chemical Bonds: 1. Chemical Bond: the force that holds 2 or more atoms together, and makes them function as a unit. a. Involve the valence electrons of elements. b. These valence electrons can be represented by using an electron dot structure (diagram). c. Electron dot diagrams are also useful in illustrating the formation of a chemical bond. 2. The strength of a bond can be determined by measuring the energy required to break that bond (called “bond energy”)

Electron Dot Diagrams • Because valence electrons are involved with chemical bonds, chemists use a short-hand method to represent atoms, called “electron dot diagrams (structures)”. • Electron dot diagram: has the following: 1. The elements symbol a. This represents the atoms nucleus and its inner-core electrons. 2. Dots surrounding the symbol, representing the atoms valence electrons. a. There will be 1 -8 dots, depending on the element

Electron Dot Diagrams • Method of creating electron dot diagram: 1. the “dots” representing the atoms valence electrons are placed one at a time on all 4 sides of the elements symbol, and only paired after all 4 sides have 1 dot. 2. It doesn’t matter where you place the initial dot.

Electron Dot Diagrams • Valence electrons: For the representative elements, the group number is also the # of valence electrons. 1. Group 1 A has 1 valence electron 2. Group 2 A has 2 valence electrons 3. Group 3 A has 3 valence electrons 4. Group 4 A has 4 valence electrons……. 5. Group 8 A has 8 valence electrons • These valence electrons are represented as the “dots” drawn around the elements symbol.

Electron Dot Diagrams • The electron dot diagrams for noble gases will show a full octet (except for helium)

Electron Dot Diagrams • Examples:

Types of Chemical Bonds • Two types of chemical bonds exist: 1. ionic bonds, forming ionic compounds. 2. Covalent bonds, forming covalent compounds. • Ionic Bonds: formed when a metal reacts with a non -metal. 1. The metal atom(s) easily give-up one or more of their valence electrons to expose a full octet of inner s & p orbitals. 2. The non-metal, having an affinity (desire) for electrons, will take on the extra electron(s), which will give it a full octet in its outer s & p orbitals.

Types of Chemical Bonds • Electron Affinity: a measure of the attraction of an atom for electrons 1. noble gases have low electron affinity ( already have a stable octet) • Reactivity of an element is directly related to its valence electrons. 1. elements will “react” in order to acquire the stable electron structure of a noble gas. 2. Forming a positive ion: atoms will lose 1+ of their valence electrons in order to obtain a noble gas configuration. 3. Forming a negative ion: atoms will gain 1+ electrons to obtain a noble gas configuration.

Covalent Bonds • There are compounds that form where both elements making up the compound need to gain electrons in order to have a stable octet. 1. They will “share” electrons to become stable. 2. Bonding follows the “Octet rule”. Octet rule: Elements will form bonds with other elements in order to obtain a stable octet of valence electrons. 3. this is called a covalent bond.

Covalent Bonds • Covalent Bond: A chemical bond that results from the sharing of valence electrons 1. Shared electrons are considered to be a part of the complete outer energy level of both atoms. 2. Generally: a. Mostly occurs in bonds between non-metals. b. Occurs with elements close to each other on the periodic table. 3. Molecule: formed when 2 or more atoms bond covalently. 4. most common materials we encounter have covalent bonds (clothing, food)

Covalent Bonds • Formation of Covalent bonds: 1. Diatomic molecules: all of them share electrons (H 2, N 2, O 2, F 2, Cl 2, Br 2, & I 2) 2. Example: Fluorine (group 7 A; 7 valence electrons) a. As 2 fluorine atoms approach each other, the protons of one will be attracted to the electrons of the other. b. They will get close enough to achieve maximum attraction (too close, and they will repel each other) c. At this point, they are sharing an electron, so both fluorine atoms have a stable octet. d. Each fluorine atom will now have 1 “bonding pair” of electrons, and 3 “lone” (unshared) pairs of electrons.

Covalent Bonds • Types of Covalent Bonds: 1. Single covalent Bonds: A covalent bond in which only a single pair of electrons is shared. a. The shared electron pair = “Bonding pair” b. Bonding pairs are represented in electron dot diagrams called Lewis structures. Lewis structure: model that uses electron-dot structures to show electrons are arranged in molecules. c. Bonding pairs are represented as a pair of dots, or a line drawn between atoms.

Covalent Bonds • Examples of single covalent bonds: 1. Hydrogen molecule (H 2) = H-H or H: H 2. All halogens (group 7 A) have 7 valence electrons, need only 1 more to form an octet, and therefore will share one pair of electrons, forming a single covalent bond. a. Can form a single covalent bond between themselves (F 2, Cl 2, I 2), and with other elements. 3. Group 6 A has 6 valence electrons, and need 2 more to form an octet. (will share 2 pairs to do it) a. Will form 2 separate single covalent bonds. b. Ex. H 2 O (on next slide)

Single covalent Bonds • Example H 2 O Lewis Structures: • A set of dots can be drawn to represent the bonding pair of electrons, or a single line.

Single Covalent Bonds • Group 5 A elements have 5 valence electrons, and need 3 more to form an octet. 1. they will “share” 3 electrons with other atoms, forming 3 single covalent bonds. 2. Example: ammonia NH 3 3. Example: nitrogen triflouride NF 3

Single Covalent Bonds • Group 4 A elements have 4 valence electrons, and need 4 more to form an octet. 1. they will “share” 4 electrons with other atoms, forming 4 single covalent bonds. 2. Example: methane CH 4 • Single covalent bonds are also called “sigma bonds” 1. symbolized by greek letter sigma σ 2. the shared electrons will be concentrated into a “bonding orbital” between the 2 atoms Bonding orbital: localized region where bonding electrons are most likely found.

Multiple covalent bonds • Multiple covalent bonds: a bond in which atoms share more than one pair of electrons. 1. Elements that commonly form multiple covalent bonds are carbon, nitrogen, oxygen & sulfur. 2. Double covalent bond: one in which 2 pairs of electrons are shared between 2 atoms. a. Two lines are drawn between 2 atoms to indicate the presence of a double covalent bond. b. Ex: oxygen molecule O 2

Multiple covalent bonds • Triple covalent bond: one in which 3 pair of electrons are being shared between 2 atoms. 1. ex: nitrogen molecule N 2 • Bond dissociation energy: 1. energy is released when a bond forms. 2. energy must be added to break a bond. Bond Dissociation energy: the amount of energy required to break a specific covalent bond. 3. triple bonds require more than double, which requires more than single covalent bonds.

Molecular structure • Molecular structure: the use of models to predict the arrangement of atoms in a molecule. 1. Structural formula: Uses elements symbols and bonds (using dots or lines) to show relative positions of atoms. 2. The structural formula can be predicted for many molecules by drawing the Lewis Structure for the molecule. 3. There are steps that should be followed to draw a Lewis Structure for a covalent compound

Steps to determine Lewis Structures • Step 1: predict the location of certain atoms. 1. Hydrogen is always a terminal (end) atom. 2. The atom with the least attraction for shared electrons in the molecule will be the central atom. a. The element is usually the one closest to the left side of the periodic table (if a choice is there). b. Central atom: located in the center of the molecule, with all other atoms being terminal atoms.

Steps to determine Lewis Structures • Step 2: Find the total number of electrons available for bonding (bonding pairs). 1. Count the total number of valence electrons of all the atoms in the molecule. 2. Divide this number by 2 to determine the number of bonding pairs. • Step 3: apply the bonding pairs of electrons in a specific way the molecule. 1. Place one bonding pair (representing a single bond) between the central atom and each of the terminal atom.

Steps to determine Lewis Structures • Step 3 (cont. ): 2. Subtract the number of pairs used from the original total number of bonding pairs. a. The remaining electron pairs will be applied as lone pairs of electrons, or used in double or triple bonds. • Step 4: Place lone pairs around each terminal atom that is bonded to the central atom 1. This is done to satisfy the octet rule for each terminal atom. 2. The exception to this is hydrogen as a terminal atom.

Steps to determine Lewis Structures • Step 5: Assign any remaining bonding pairs of electrons to the central atom • Step 6: Evaluate if the central atom has an octet around it (4 electron pairs). 1. If central atom has an octet, and all bonding pairs have been used, you are done! 2. If central atom doesn’t have an octet, then you must convert one or more of the lone pairs of electrons on the terminal atom(s) to a double or triple bond between that terminal atom and the central atom. a. Often, carbon, nitrogen, oxygen or sulfur are the central atoms (for double & triple bonds)

Steps to determine Lewis Structures • Additional step: if Lewis structure being drawn is for a polyatomic ion. 1. The atoms within the polyatomic ion are covalently bonded to each other. 2. When determining the number of bonding pairs of electrons that exist for the polyatomic ion, you must take into account the charge of the ion. 3. Find the total number of valence electrons of all involved atoms, then add/subtract electrons to match the charge of the polyatomic ion. a. Apply the electrons pairs as previously described.

Example of a Lewis Structure for a Polyatomic Ion • Chlorite ion: Cl. O 2 -

Lewis structures with 2 central atoms • Some molecules have 2 central atoms that are bonded to each other, and have terminal atoms on each of them. • Ex: C 2 H 4 (ethene) 1. follow same steps to form the Lewis structure • Exceptions to octet rule: 1. Incomplete octets 2. Expanded octets a. always apply extra pairs to central atom

Exceptions to Octet Rule • Some molecules do not obey the octet rule. • 3 Reasons: 1. There a small group of molecules that have an odd number of valence electrons, and can’t form an octet around each atom: a. Ex: NO 2 N = 5 valence electrons O = 6 valence electrons *5+6+6 = 17 valence electrons, therefore you can’t make exact pairs (one left over) b. CLO 2 - and NO are additional examples

Exceptions to Octet Rule • Reason # 2: certain compounds will form with fewer than 8 electrons present around its central atom (rare) 1. Ex: BH 3 • Reason # 3: certain compounds have central atoms that have more than 8 valence electrons, called an “expanded octet”. 1. Ex: PCl 5 • Any lone pairs not used should be applied to central atom.

Resonance Structures • When there are more than one correct Lewis structures that can be drawn for a molecule, that molecule is said to “show resonance”. Resonance: a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion. 1. Often seen when a molecule has a double or triple bond. 2. The two or more correct Lewis structures representing a single molecule/ion are called “resonance structures”.

Resonance Structures • When drawing resonance structures, the position of the atoms will remain the same. 1. the difference between the resonance structures is the position of the lone and bonding pairs of electrons 2. Ex: NO 33. Ex: CO 2

Molecular Shape • General info: 1. The shape of a molecule determines its physical & chemical properties. 2. The shape of a molecule is determined by the overlap of orbitals that share electrons. 3. Many chemical reactions depend on the ability of 2 compounds to contact each other. a. The shape of molecules dictates how close they will get to one another 4. A Lewis structure must be drawn first in order to determine molecular shape. 5. The model used to determine molecular shape is Valence Shell Electron Pair Repulsion Model (VSEPR)

VSEPR Model • Model is based on an electron arrangement that minimizes the repulsion of shared/unshared pairs of electrons around the central atom. 1. This repulsion results in atoms being oriented at fixed angles to each other. 2. The angle formed by any 2 terminal atoms and the central atom = Bond Angle • Shared pairs of electrons occupy a bonding orbital which is smaller • Lone pairs of electrons are not shared by 2 different nuclei, and therefore occupy a larger orbital 1. They will push the shared bonding orbitals closer together (and give shape to the molecule)

VSEPR Model • Molecules will have different geometric shapes, depending on: 1. The number of bonding pairs of electrons between the central atom & terminal atoms. 2. The presence or absence of lone pairs of electrons on the central atom. • Molecular Shapes: 1. Linear 2. Trigonal Planar 3. Tetrahedral 4. Trigonal Pyramidal 5. Bent

Molecular Shape • Molecules containing NO lone pairs of electrons on the central atom: 1. Ex: Beryllium Chloride 2. Bonding electrons will seek maximum separation. 3. Molecular shape will be Linear 4. Bond Angle: 180° 5. Has 2 pairs of electrons shared off the central atom. 6. Doesn’t exhibit polarity (non-polar molecule)

Molecular Shape • Molecules containing NO lone pairs of electrons on the central atom: 1. Ex: Carbon Dioxide 2. Bonding electrons will seek maximum separation. 3. Molecular shape will be Linear 4. Bond Angle: 180° 5. Has 4 pairs of electrons shared off the central atom 6. Doesn’t exhibit polarity (non-polar molecule)

Molecular Shape • Molecules containing NO lone pairs of electrons on the central atom: 1. Ex: Aluminum Chloride 2. Bonding electrons will seek maximum separation. 3. Molecular shape will be Trigonal Planar 4. Bond Angle: 120° 5. Has 3 pairs of electrons shared off the central atom 6. Doesn’t exhibit polarity (non-polar molecule)

Molecular Shape • Molecules containing NO lone pairs of electrons on the central atom: 1. Ex: methane 2. Bonding electrons will seek maximum separation. 3. Molecular shape will be Tetrahedral 4. Bond Angle: 109. 5° 5. Has 4 pairs of electrons shared off the central atom. 6. Doesn’t exhibit polarity (non-polar molecule)

Molecular Shape • Molecules containing ONE lone pair of electrons on the central atom: 1. Ex: PH 3 2. Lone pair of electrons takes up greater space, forcing the 3 single covalent bonds away. 3. Molecular shape will be Trigonal pyramidal 4. Bond Angle: 107. 3° 5. Has 3 pairs of electrons shared off the central atom, and 1 lone pair. 6. Exhibits polarity (polar molecule)

Molecular Shape • Molecules containing TWO lone pairs of electrons on the central atom: 1. Ex: Water 2. Two lone pairs take up greater space, forcing the 2 single covalent bonds away. 3. Molecular shape will be Bent 4. Bond Angle: 104. 5° 5. Has 2 pairs of electrons shared off the central atom, and 2 lone pairs. a. Lone pairs will both be on one side. 6. Exhibits polarity (polar molecule)

Electronegativity Difference & Bond Character • The type of chemical bond that will form between 2 elements can be predicted by determining the electronegativity difference of the 2 atoms that are bonded. • When covalent bonds form between atoms of different elements, there will be an unequal sharing of their electron pairs, due to differences in electronegativity values. 1. This unequal sharing of the bonding pairs of electrons results in a Polar Covalent Bond

Electronegativity Difference & Bond Character • When covalent bonds form between atoms of different elements, large differences in electronegativity values between the atoms will result in an ionic bond. 1. Bonding is therefore often not clearly ionic or covalent. 2. As the difference in electronegativity increases, the bond will become more ionic. 3. With atoms of the same type (diatomic), the electronegativity difference is zero, resulting in a Pure Covalent Bond. 4. Smaller electronegativity differences will result in a Polar Covalent Bond.

Electronegativity Difference & Bond Character

Electronegativity Difference & Bond Character • Electronegativity difference of zero = pure covalent • Electronegativity difference of 1. 70 = 50% covalent 50% ionic • Electronegativity difference above 1. 70 = more ionic

Polar Covalent Bonds • Between pure covalent bonds and ionic bonds, there is a range of bonds that are polar covalent. 1. One of the atoms in the bond are more electronegative than the other, causing an un-equal sharing of the bonding electron pair (like un-equal sides in a tug-of-war contest). 2. These shared electrons will be pulled closer to the more electronegative atom, causing a “polarity” to the molecule (one side of molecule will be slightly positive, the other side slightly negative. 3. The delta symbol “δ” is used with a +/- sign drawn next to the molecule to show which side is slightly positive, and which is slightly negative.

Polar Covalent Bonds • The result of a polar covalent bond: a Dipole (twopole) molecule. 1. one side will be slightly negative δ-, and the other side slightly positive δ+. 2. Example: H 2 O 3. polar molecules will be attracted by an electrical field. 4. Some molecules exhibit polarity between their central & terminal atoms, but their overall polarity is balanced, so that the molecule is considered to be non-polar.

Polar vs Non-polar molecules • Polar molecules: 1. Attracted by an electrical field 2. Have lone pairs of electrons, which push away bonding pairs. 3. Have slightly positive and slightly negative sides. 4. Are asymmetrical in shape. • Non-polar molecules: 1. Are not attracted by an electrical field 2. Do not have any lone pairs of electrons. 3. Are symmetrical in shape.