Chapter 8 and 9 notes Chemical Bonds Chemical

Chapter 8 and 9 notes

Chemical Bonds – Chemical bond – the force that holds two atoms together • Chemical bonds may form by the attraction between a positive nucleus and negative electrons or the attraction between a positive ion and a negative ion

– The elements within a group on the periodic table have similar properties. Many of these are due to the number of valence electrons • These same electrons are involved in the formation of chemical bonds between two atoms • Recall that an electron dot structure is a type of diagram used to keep track of valence electrons and is especially when illustrating the formation of chemical bonds

• Recall from Chapter 6 that ionization energy refers to how easily an atom loses an electron • The electron affinity indicates how much attraction an atom has for electrons. • Nobel gases, have high ionization energies and low electron affinities, show a lack of chemical reactivity – the difference in reactivity is directly related to the valence electrons

– All atoms have valence electrons. Noble gases have electron configurations that have a full outermost energy level. • The presence of eight valence electrons in the outer energy level is chemically stable and is called a stable octet • Elements tend to react to acquire the stable electron structure of a noble gas

8. 2 The formation and nature of ionic bonds Objectives 1. Describe the formation of ionic bonds 2. Account for many of the physical properties of an ionic compound • 3. Discuss the energy involved in the formation of an ionic bond. • •

• The electrostatic force that holds oppositely charged particles together in an ionic compound is an ionic bond • Compounds that contain ionic bonds are ionic compounds • ionic bond occurs between metals and nonmetal, it is a transfer of electrons

• Hundreds of compounds contain ionic bonds • Many ionic compounds are binary which means they contain only two different elements • Binary ionic compounds contain a metallic cation and a nonmetallic anion.

Chemical Bonding Ionic Bonds: atoms give up or gain electrons and are attracted to each other by coulombic attraction Na loses an e– Na 1+ + e– ionic compounds = salts Cl gains an e– Cl + e– Cl 1– Na 1+ + Cl 1– Na. Cl K 1+ + NO 31– KNO 3 ***where NO 31– is a polyatomic ion: a charged group of atoms that stay together ***(polyatomic ion = atoms bonded together that exhibit an overall charge)

Ionic Bonding Na. Cl n=3 - n=2 - - - - Na [Ne]3 s 1 - - - + - - - - - Cl [Ne]3 s 23 p 5 - - - Na+ [Ne] - - - Cl[Ne]3 s 23 p 6 Transfer of electrons to achieve a stable octet (8 electrons in valence shell).

• Properties of Ionic Compounds • A. The chemical bonds that occur between the atoms in a compound determine many of the physical properties of the compound • During the formation of ionic compounds, the positive and negative ions are packed into a regular repeating pattern that balances the forces of attraction and repulsion between the ions. This packing forms an ionic crystal.

• Large numbers of positive ions and negative ions exist together in a ratio determined by the number of electrons transferred from the metal to the nonmetal • The strong attraction of positive ions and negative ions in an ionic compound results in a crystal lattice. A crystal lattice is a three dimensional geometric arrangement of particles. In the lattice, each positive ion is surrounded by negative ions and each negative ion is surrounded by positive ions. Ion crystals vary in shape due to the size and relative numbers of the ions bonded

Types of Bonds Ionic Bonding - Crystal Lattice Table salt

• Melting point, boiling point, and hardness are physical properties that depend on how strongly the particles are attracted to each other. • Because ionic bonds are relatively strong, the crystals that result require a large amount of energy to be broken apart. • Therefore, ionic crystals have high melting points and boiling points

• the formation of ionic compounds from positive and negative ions is always exothermic • The attraction of the positive ion for the negative ions close to it forms a more stable system that is lower in energy than the individual ions.

– The energy required to separate one mole of the ions of an ionic compound is referred to as the lattice energy. The strength of the forces holding ions in place is reflected by the lattice energy. The more negative the lattice energy the stronger the force of attraction – Lattice energy is directly related to the size of the ions bonded. The value is also affected by the charge of the ion.

• How to write compounds! Criss- Cross 1. Write the element with its charge 2. Criss cross the numbers down to the other element (leaving the charges off)

Criss-Cross Rule Example: Aluminum Chloride Step 1: Aluminum Chloride Step 2: Al 3+ Cl 1 - Step 3: Al 1 Cl 3 Step 4: Al. Cl 3

Criss-Cross Rule Example: Aluminum Oxide Step 1: Aluminum Oxide Step 2: Al 3+ O 2 - Step 3: Al 2 O 3 Step 4: Al 2 O 3

Criss-Cross Rule Example: Magnesium Oxide Step 1: Magnesium Step 2: Mg 2+ O 2 - Step 3: Mg 2 O 2 Step 4: Mg 2 O 2 Step 5: Mg. O Oxide

8. 4 • 8. 4 Metallic Bonds and the Properties of Metals • Objectives: 1. describe a metallic bond • 2. Explain the physical properties of metals in terms of metallic bonds • 3. Define and describe alloys

• Metallic Bonds • Although metals are not ionic, they share several properties with ionic compounds, • they form lattices in the solid state which are similar to ionic crystal lattices. In such a lattice, 8 -12 other metal atoms surround each metal atom

• Although metal atoms have at least one valence electron they do not share these electrons with neighboring atoms, nor do they lose electrons to form ions • Instead the energy levels of the atoms overlap. • The electron sea model proposes that all the metal atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons. • The electrons present in the outer energy levels of the bonding metallic atom are not held in place by any specific atom and can move easily from one atom to the next

• Because they are free to move they are often called delocalized electrons • When the atom’s outer electrons move freely throughout the solid a metallic cation is formed. Each such ion is bonded to all neighboring metal cations by the sea of valence electrons. • A metallic bond is the attraction of a metallic cation for delocalized electrons

• Properties of metals – the typical physical properties of metals can be explained by metallic bonding. These properties provide evidence of the strength of metallic bonds • in general, metals have moderately high melting points and high boiling points. The melting points are not as extreme It does not take an extreme amount of energy for them to be able to move past each other. However, during boiling, atoms must be separated from the group of cations and electrons which requires more energy.

• Metals are malleable, which means they can be hammered into sheets and ductile which means they can be drawn into wire. • Metals are generally durable. And although metallic cations are mobile in metal they are strongly attracted to the electrons surrounding them and aren’t easily removed from the metal • Delocalized electrons in a metal are free to move, keeping metallic bonds in tact. The movement explains why metals are good conductors. • As the number of delocalized electrons increases, so do the properties of hardness and strength

Chapter 9 • 9. 1 The Covalent bond • Objectives: 1. Apply the octet rule to atoms that bond covalently • 2. Describe the formation of single, double, and triple covalent bonds • 3. Compare and contrast sigma and pi bonds • 4. Relate the strength of covalent bonds to bond length and bond dissociation energy

• Why do atoms bond. • You learned in chapter 6 that all noble gases have particularly stable electron arrangements. • This is because it has a full outer energy level. • A full outer energy level consists of two valence electrons for helium and eight valance electrons for all other noble gases. • You also learned in chapter 8 that when metals and nonmetals react to form binary ionic compounds, electrons are transferred, and the resulting ions have noble gas electron configurations.

• But sometimes two atoms that both need to gain valence electrons to become stable have a similar attraction for electrons • Sharing of electrons is another way that these atoms can acquire the electron configuration of noble gases. • Remember Chapter 6 says the octet rule states that atoms lose, gain or share electrons to achieve a stable configuration of eight valence electrons, or an octet. There are some exceptions but this is the basic rule

• What is a covalent bond? • The chemical bond that results from the sharing of valence electrons is a covalent bond. • In covalent bonding the shared electrons are considered to be part of the complete outer energy level of both atoms involved. • Covalent bonding generally occurs when elements are relatively close to each other on the periodic table.

Chemical Bonding Covalent Bonds …atoms share electrons to get a full valence shell C 1 s 2 2 p 2 (4 v. e–) F 1 s 2 2 p 5 (7 v. e–) both need 8 valence electrons for a full outer shell (octet rule)

Covalent Bonding n=2 - - n=1 - - - + - - - - O [He]2 s 22 p 4 O 2 Sharing of electrons to achieve a stable octet (8 electrons in valence shell).

• The majority of covalent bonds form between nonmetallic elements • ***A molecule is formed when two or more atoms bond covalently and produce a neutral particle. • The sugars you eat; the proteins, fats, and DNA in your body; and the wool, cotton and fibers in your clothes all consist of molecules formed from covalently bonded atoms.

• Formation of a covalent bond • Hydrogen H 2, nitrogen N 2, oxygen O 2, fluorine F 2, chlorine Cl 2, bromine Br 2 and iodine occur in nature as diatomic molecules, not as single atoms because the molecules formed are more stable that the individual atoms. • 7 diatomic elements: • ***These form pure covalent bonds or non polar molecules! • H, N, F, O, I, Cl, Br

• the diatomic molecule exists because the sharing of one pair of electrons will give both fluorine atoms stable noble gas configurations. • Each fluorine atom in the fluorine molecule has one shared pair of electrons and three lone pairs, which are unshared pairs of electrons.

Covalent bonding • • Fluorine has seven valence electrons A second F atom also has seven By sharing electrons Both end with full orbitals (stable octets) F 8 Valence electrons

• Each covalently bonded atom equally attracts one pair of shared electrons, thus, two electrons shared by two hydrogen nuclei belong to each atom simultaneously. And both Hydrogens have the noble gas configuration. • When a single pair of electrons is shared, such as in hydrogen a single covalent bond forms. The shared electron pair often referred to as the bonding pair, is represented by either a pair of dots or a line in the Lewis structure. • Lewis Structures use electron dot diagrams to show electrons are arranged in molecules

Lewis Structure Lewis structure: a model of a covalent molecule that shows all of the valence electrons 1. ***Two shared electrons make a single covalent bond, four make a double bond, etc. 2. unshared pairs: pairs of un-bonded valence electrons 3. Each atom needs a full outer shell, i. e. , 8 electrons. Exception: H needs 2 electrons

Lewis Structure carbon tetrafluoride (CF 4) x x x o o Co o x x F x x x x F x o x x x x o x x F x x Co F ox x x x x C x x F x covalent compounds = molecular compounds (have lower melting points than do ionic compounds) x x F x x x x

Lewis Structure o o methane (CH 4) o o N o o C o I x x x Ox x x Ox o o o x x o I N I x x ox x x I x x x o o H Hx C ox H x x o o x o H o o nitrogen triiodide (NI 3) carbon dioxide (CO 2) o C o H H C H H H x x o o C o x x x x Ox x x I N x. Ix x x x I x xx x x o o xx O =C=O xx xx

• Multiple covalent bonds • In many molecules, atoms attain a noble-gas configuration by sharing more than one pair of electrons between two atoms, forming a multiple covalent bond.

• Strength of covalent Bonds • A covalent bond involves attractive and repulsive forces. • Several factors control the strength of covalent bonds – The strength depends on how much distance separates bonded nuclei • The distance between the two bonding nuclei at the position of maximum attraction is called bond length, which is determined by the size of the atoms and how many electron pairs are shared

• **As the number of shared electron pairs increases, bond length decreases. The shorter the bond length the stronger the bond. • An energy change accompanies the forming or breaking of a bond between atoms in a molecule.

• **Energy is released when a bond forms and energy must be added to break the bonds in a molecule. • The energy amount required to break a specific covalent bond is called bond dissociation energy. • **Breaking bonds always requires the addition of energy

Types of formulas and bonds • Chemical formula: shorthand representation of a compound using atomic symbols and subscripts. • Example: Na. Cl _____(____ atoms) • Ca 3(PO 4) _____(____ atoms) • (NH 4)2 SO 4 _____(____ atoms)

Types of formulas and bonds • Chemical formula: shorthand representation of a compound using atomic symbols and subscripts. • Example: Na. Cl sodium chloride(2 atoms) • Ca 3(PO 4)2: calcium phosphate) (13 atoms) • (NH 4)2 SO 4 Amonium sulfate(15 atoms)

Molecular formula • Shows the kind of atom and the number of atoms in a molecule. Nonmetals only. These may look like they can be reduced • Examples • H 2 O _______(____ atoms) • C 6 H 6 _______(____ atoms) • C 12 H 22 O 11 _______(____ atoms)

Molecular formula • Shows the kind of atom and the number of atoms in a molecule. Nonmetals only. These may look like they can be reduced • Examples • H 2 O water (3 atoms) • C 6 H 6 Benzene (12 atoms) • C 12 H 22 O 11 sucrose (45 atoms)

Structural formula • Indicates the kind of atom, number of atoms and the arrangement of bonds of the atoms. This gives the most information • CH 4 – a single dash is a single bond • C 2 H 4 – two dashes = double bond • C 2 H 2 – three dashed = triple bond • Compounds of only carbon and hydrogen are hydrocarbons

9. 3 • 9. 3 Molecular Structures • Objectives • List five basic steps used in drawing Lewis structures • Explain three exceptions to the octet rule, and identify molecules in which these exceptions occur

• To Determine the shape of the atom: • The following procedures should be followed to determine Lewis structures • 1. Predict the location of certain atoms • Hydrogen is always a terminal, or end atom. Because it can share only one pair of electrons, hydrogen can be connected to one other atom • The atom with the least attraction for shared electrons in the molecule is the central atom. This element usually is the one closer to the left on the periodic table. The central atom is located in the center of the molecule, and all other atoms become terminal atoms

• 2. Find the total number of electrons available for bonding. This total is the number of valence electrons in the atoms in the molecule • 3. Determine the number of bonding pairs by dividing the number of electrons available for bonding by two • 4. Place one bonding pair (single bond) between the central atom and each of the terminal atoms

• 5. Subtract the number of pairs you used in step 4 from the number of bonding pairs you determined in step 3. • The remaining electron pairs include lone pairs as well as pairs used in double and triple bonds. • Place lone pairs around each terminal atom bonded to the central atom to satisfy the octet rule. Any remaining pairs are assigned to the central atom

• 6. If the central atom is not surrounded by four electron pairs, It does not have an octet. • You must convert one or two of those lone pairs on the terminal atoms to a double bond or a triple bond between the terminal atom and the central atom. • These pairs are still associated with the terminal atom as well as the central atom. • Remember that in general, carbon, nitrogen, oxygen and sulfur can form double or triple bonds with the same element or with another element.

• The process for drawing Lewis structures for polyatomic ions is similar to drawing them for covalent compounds. • The main difference is in finding the total number of electrons available for bonding. Compared to the number of valence electrons present in the atoms that make up the ion, more electrons are present if the ion is negatively charged and fewer are present if the ion is positive. • To find the total number of electrons available for bonding, first find the number available in the atoms present ion • Then subtract the ion charge if it is positive, and add the charge if it is negative

• Exceptions to the Octet rule • The Lewis structure is focused on the attainment of an octet by all atoms when they bond with other elements. Some do not obey the octet rule – Hydrogen wants 1 pair – Be wants 2 pairs – B and Al want 3 Pairs • Three reasons exist for these exceptions

• First a small group of molecules has an odd number of valence electrons and cannot form an octet around each atom • Ex, NO 2 has 5 valence electrons from nitrogen and 12 from oxygen = 17 electrons.

• Second, some compounds form with fewer than eight electrons around an atom. BH 3

• A total of six electrons are shared by the boron atom. • Such compounds tend to be reactive and can share an entire pair of electrons donated by another atom. • One atom donates a pair of electrons to be shared with an atom or ion that needs two electrons to become stable, a coordinate covalent bond forms.

• The third group that doesn’t follow the octet rule has central atoms that contain more than eight valence electrons. • This electron arrangement is referred to as an expanded octet. • It can be explained by considering the d orbital that occurs in the energy levels of elements in period three or higher • An example is the molecule PCl 5. Five bonds are formed with ten electrons shared in one s orbital, three p orbitals and one d orbital

Practice • • • NH 3 Si. F 4 H 2 S As. Cl 3 CH 3 Br

Practice • NH 3 • Si. F 4 • H 2 S • As. Cl 3 • CH 3 Br

CO 3 -2 has three possible lewis structures. They are all possible, this is an exapleof resonance (when a molecule cannot be correctly represented by a single lewis structure) Double arrows are drawn between equivalent structures. Note that a charge of 2 must be added to the final structure.

9. 4 • 9. 4 Molecular shapes • Objectives – Discuss the VSEPR bonding theory – Predict the shape of and the bond angels in a molecule – Define Hybridization

• The shape of a molecule determines many of its physical and chemical properties – Molecular shape in turn is determined by the overlap of orbitals that share electrons. • VSEPR Model – Many chemical reactions depend on the ability of two compounds to contact each other

• The shape of the molecule determines whether or not molecules can get close enough to react • Once a Lewis structure is drawn, you can determine the molecular geometry or shape of the molecule • ***The model used to determine the molecular shape is referred to as the Valence Shell Electron Pair Repulsion Model or VSEPR model. This model is based on an arrangement that minimizes the repulsion of shared and unshared pairs of electrons around a central atom.

• ***The repulsions among electron pairs in a molecule result in atoms existing at fixed angles to each other. • The angle formed by any two terminal atoms and the central atom is a bond angle. • Bond angles predicted by VSEPR are supported by experimental evidence • Shared electron pairs repel one another. • Lone pairs occupy a slightly larger orbital because they are not shared • Shared bonding orbitals are pushed together slightly by lone pairs

• Hybridization • A hybrid results from combining two of the same type of object, and it has characteristics of both. • Hybridization is a process in which atomic orbitals are mixed to form new, identical hybrid orbitals

• Electronegativity and Polarity • Objectives: • Describe how electronegativity is used to determine bond type • Compare and contrast polar and nonpolar covalent bonds and polar and nonpolar molecules • Describe the characteristics of compounds that are covalently bonded

• Electronegativity Difference and Bond Character • Electron affinity is a measure of the tendency of an atom to accept an electron • Excluding Noble gases, electron affinity increases as the atomic number increases within a given period and decreases with an increase in atomic number within a group. • The scale of electronegtivities allows a chemist to evaluate the electron affinity of specific atoms when they are incorporated into a compound. • Fluorine has the highest electronegativity value and francium has the lowest.

Electronegativity Guidelines • • If > 2. 0 the bond is Ionic If < 1. 7 the bond is Covalent > 0. 3 the bond is Polar Covalent < 0. 3 the bond is Non-Polar • If the difference is between 1. 7 and 2. 0 and involves a metal the bond is considered ionic. If only nonmetals are involved the bond is polar covalent. Example HF is polar covalent Example HCl Electronegativity Cl = 3. 16 Electronegativity H = 2. 20 Difference = 0. 96 Polar Covalent

Try these • • • Cesium and fluorine _______ Fluroine and silicon _______ Hydrogen and chlorine _______ Magnesium and nitrogen _______ Beryllium and fluorine _______ Chlorine and bromine _______ Chlorine and lithium _______ Hydrogen and iodine _______ Strontium and oxygen _______

Try these • • • Cesium and fluorine 4 -0. 79=3. 2 ionic Fluroine and silicon - ionic Hydrogen and chlorine- polar covalent Magnesium and nitrogen- ionic Beryllium and fluorine - ionic Chlorine and bromine – non polar covalent Chlorine and lithium – ionic Hydrogen and iodine – polar covalent Strontium and oxygen - ionic

• Two atoms will bond if it lowers their potential energy. Nature seeks a decrease in energy • Memorize • **Breaking a chemical bond usually absorbs energy (Endothermic) • Molecular substances have covalent bonds • Diatomic molecules have nonpolar covalent molecules. H N F O I Cl Br • Polyatomic ions – a group of charged atoms that are covalently bonded (all nonmetals) • Compounds composed of only carbon and hydrogen are called hydrocarbons

• The character and type of chemical bond can be predicted using the electronegativity difference of the elements that are bonded. • For identical atoms which have an electronegativity difference of zero, the electrons in the bond are equally shared between the two atoms and the bond is considered ***nonpolar covalent, which is a*** pure covalent bond. **Example diatomic molecules • Chemical bonds between atoms of different elements are never completely ionic or covalent and the character of the bond depends on how strongly bonded atoms are attracted. • ***Unequal sharing of electrons results in a polar covalent bond.

• Large differences in electronegativity indicate a bond is primarily ionic • As the difference in electronegtivity increases the bond becomes more ionic in character.

• Polar covalent Bonds • Why are some bonds polar covalent? Sharing is not always equal • Polar covalent bonds form because not all atoms that share electrons attract them equally

• When a polar bond forms, the shared pair of electrons is pulled toward one of the atoms. • The electron spends more time around that atom than they do the other. • Partial charges occur at the ends of the bonding. • The more electronegative atom is located at the partially negative end while the less electronegative atom is found at the partially positive end. • The resulting polar bond often is referred to as a dipole (two poles)

• A polar molecule has a partial positive and partial negative side, make it a dipole because of the two partial charges

• Polar Molecule or not • How do you tell if it is polar or not. Let’s look at H 2 O and CCl 4. The electronegativity differences are 1. 24 and 0. 61. Both are polar covalent

• Both molecules contain more than one polar covalent bond. But water molecules are more polar. Let’s look at the geometry

**The shape of a molecule as well as the polarity of its bonds, determines the polarity of the molecule

• The shape of H 2 O determined by VSEPR is bent because of the two lone pairs of electrons. Because it is not symmetric the molecule has a definite positive end and negative end, thus it is polar • The shape of CCl 4 is tetrahedral and symmetric. The electrical charge is identical throughout the molecule. This makes it nonpolar. • Symmetric molecules are usually nonpolar and asymmetric are polar. (Remember: SNAP) • **The POLARITY of a large molecule helps determine its shape

• Properties of Covalent Compounds • Differences in properties are a result of differences in attractive forces • In covalent compounds the covalent bond between atoms is strong, but attraction between individual molecules is relatively weak