1 Chemical Bonds Elements form bonds to be in a lower energy state 1. Ionic Bonds transfer of electrons, between metal and nonmetal 2. Covalent Bonds sharing of electrons, between two nonmetals 3. Metallic Bonds- neighboring atoms in solid metals form bonds Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons to achieve a stable octet (noble gas configuration)
2 Electron Dot Symbols Valence electrons: reside in the highest occupied energy level, reside in the outer s & p orbitals and are the electrons involved in chemical bonding. Electron-dot symbols are convenient way of showing the s & p electrons & tracking them in bond formation -They consist of the chemical symbol for the element plus a dot for each valence electron Consider sulfur whose electron configuration is [Ne]3s 2 3p 4, thus there are six valence electrons: S
3 Ch. 7 Ionic and Metallic Bonding I II III IV
4 B. Ionic Bonds Ionic Bonds atoms transfer electrons from a cation (positive ion) to an anion (negative ion) to achieve an octet. Ionic compounds are stable due to the electrostatic forces between unlike charges organizing the ions of ionic substances into a rigid, organized three-dimensional arrangement: The ions are drawn together Energy is released Ions form solid lattice structure
5 Lattice Energy Lattice energy: The energy required to completely separate a given quantity of a solid ionic compound into its gaseous ions. Thus, in reverse, the high energy is given off as heat and light when Na + and Cl - is incorporated into the NaCl salt lattice.
6 Steps in Ionic Bonding Process (1) Ionization Energy (IE) Step 1: The minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion. First ionization energy: Na(g) + IE 1 Na + (g) + e - ; + H (positive/ endothermic) Second Ionization Energy: Be(g) + IE 1 + IE 2 Be 2+ (g) + 2 e - ; + H (positive) The greater the IE, the more difficult it is to remove an e -.
7 Steps in Ionic Bonding Process (2) Electron Affinity (EA): Most atoms can attract e - to form negatively charged ions The energy change that occurs when an e - is added to a gaseous atom. For most atoms, the energy released when an e - is added. Cl(g) + e - Cl - (g) + electron affinity (negative/exothermic) ; - H The greater the attraction between a given atom and an added e -, the more negative the atom s EA. Halogens s 2 p 5 have the most negative EA.
8 Steps in Ionic Bonding Process (3) Lattice Energy (LE) The release of energy that occurs when ions of opposite charge are attracted to each other and form a stable ionic compound. Na + + Cl - NaCl + lattice energy ; - H (negative/exothermic) Ionic compounds have very large LE which makes up for endothermic ionization energy.
9 B. Properties of Ionic Compounds Most ionic compounds are crystalline solids at room temperature Arranged in repeating threedimensional patterns Ionic compounds generally have high melting points Large attractive forces result in very stable structures
10 B. Properties of Ionic Compounds Ionic compounds can conduct an electric current when melted or dissolved in water When ionic compounds are dissolved in water the crystalline structure breaks down so ions are able to move freely which results in conductivity
11 The positive Na ions move to the cathode and the negative Cl ions move to the anode
12 Ch. 7 Ionic and Metallic Bonding III. Bonding in Metals (p ) I II III IV
13 A. Metallic Bonding Metallic bonds: Consist entirely of metal atoms. Bonding is due to valence electrons which are delocalized throughout the entire solid The metal is held together by the strong forces of attraction between the positive nuclei and the delocalized electrons.
14 B. Metals Metals are good conductors of heat and electricity because the valence electrons are able to flow freely Valence electrons of metals can be thought of as a sea of electrons, very mobile
15 C. Metallic Bond Metallic Bonding - Electron Sea
16 D. Metallic Properties Have luster, are ductile and malleable Luster = shine Ductile = ability to be drawn into wires Malleable = ability to be shaped, pounded, etc
17 D. Metallic Properties Properties can be explained by the mobility of electrons in metals When subjected to pressure, the cations easily slide past each other like a ball bearing immersed in oil.
18 B. Types of Bonds Bond Formation Smallest Unit Physical RT Melting Point Solubility in Water Electrical Conductivity Other Properties METALLIC e - are delocalized among metal atoms electron sea solid very high no yes (any form) malleable, ductile, lustrous
19 Covalent Bonding
20 What is a covalent bond? A chemical bond that results from the sharing of electrons, to form a stable octet or duet (Hydrogen only needs 2 to be stable) Molecule = two or more atoms that are held together by covalent bonds H 2 O Majority of covalent bonds form between nonmetals (CLOSE together on periodic table)
22 Covalent Bonding Formation Diatomic molecule molecule containing the same two atoms Some elements always exist this way because they are more stable than the individual atoms Cl 2
23 B. Diatomic Elements The Seven Diatomic Elements Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 H N O F Cl Br I
24 Bonds in all the polyatomic ions and diatomics are all covalent bonds
25 Single Covalent Bonds Two atoms held together by a sharing of one pair of electrons are joined together by a single covalent bond.
26 Single Covalent Bonds An electron dot structure represents the shared pair of electrons of the covalent bond by two dots. A structural formula represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms
27 Single Covalent Bonds A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair of a nonbonding pair. Lone pair
28 Double and Triple Covalent Bonds Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two or three pairs of electrons. A double bond involves sharing two pairs of electrons. A triple bond involves sharing three pairs of electrons.
29 Bond Length From a study of various Nitrogen containing compounds bond distance as a function of bond type can be summarized as follows: N-N 147pm N=N 124pm N N 110pm As a general rule, the more e - that are shared: the stronger the covalent bond (N N > O=O > F F) the shorter the covalent bond (N N < O=O < F F)
30 Double and Triple Covalent Bonds
31 Molecular Structure Lewis Diagrams (p ) I II III
32 Drawing Lewis Diagrams 1. Arrange atoms Singular atom is usually in the center (often Carbon) If not Carbon, least e - neg atom is in center Hydrogen is always terminal 2. Find total # of e - available to bond (valence e - ) 3. Place a pair of electrons between central atom and each terminal atom
33 Drawing Lewis Diagrams 4. Place remaining electrons in pairs around terminal atoms (except H) to satisfy octet rule Any remaining pairs are assigned to central atom 5. Determine whether or not central atom satisfies octet If not, convert one or more lone pairs from terminal atoms to double or triple bonds Certain atoms can be exceptions to octet rule H, Be, B, S, P, Xe
34 Drawing Lewis Diagrams CF 4 1 C 4e - = 4e - 4 F 7e - = 28e - 32e - - 8e - 24e - F F C F F
35 Drawing Lewis Diagrams CO 2 1 C 4e - = 4e - 2 O 6e - = 12e - 16e - - 4e - 12e - O C O
36 Polyatomic Ions To find total # of valence e - : Add 1e - for each negative charge Subtract 1e - for each positive charge Place brackets around the ion and label the charge
37 Polyatomic Ions ClO Cl 7e - = 7e - 4 O 6e - = 24e - 31e - + 1e - 32e - - 8e - 24e - O O Cl O O
38 Octet Rule F Exceptions: Hydrogen 2 valence e - Boron & Beryllium get 6 & 4 valence e - respectively B F H F N O S O H F F F F F Very unstable!! Expanded octet more than 8 valence e - (e.g. S, P, Xe)
39 Drawing Lewis Diagrams BeCl 2 1 Be 2e - = 2e - 2 Cl 7e - = 14e - 16e - - 4e - 12e - Cl Be Cl
40 Drawing Lewis Diagrams SF 6 1S 6e - = 6e - + 6F 7e - = 42e - 48e - F F F S F 48 e e - 36 e - F F
41 Resonance Structures Molecules that can t be correctly represented by a single Lewis diagram Actual structure is an average of all the possibilities Show all possible structures separated by double-headed arrows
42 C. Resonance Structures SO 3 O O S O O O S O O O S O
43 Bond Polarity Most bonds are a blend of ionic and covalent characteristics. Difference in electronegativity determines bond type.
44 Bond Polarity Electronegativity Attraction an atom has for a shared pair of electrons. higher e - neg atom - lower e - neg atom +
45 Electronegativity Difference If the difference in electronegativities is between: 1.7 to 4.0: Ionic Greater than 0.3 & less than 1.7: Polar Covalent 0.0 to 0.3: Non-Polar Covalent
46 The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are bonded. Compound F 2 or F-F CF 4 LiF or Li-F Electronegativity Difference Type of Bond = = = 3.0 Non-polar covalent (strong) Polar covalent Ionic (noncovalent)
47 Bond Polarity Nonpolar Covalent Bond e - are shared equally symmetrical e - density usually identical atoms Ex: H 2 or Cl 2
48 Bond Polarity Polar Covalent Bond e - are shared unequally asymmetrical e - density results in partial charges (dipole) Ex: H 2 O + -
49 Polar Covalent Bonds: Unevenly matched, but willing to share.
50 - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.
51 Ch. 8 Molecular Structure Molecular Geometry (p ) I II III
52 VSEPR Theory Valence Shell Electron Pair Repulsion Theory Electron pairs orient themselves in order to minimize repulsive forces
53 VSEPR Theory Types of e - Pairs Bonding pairs form bonds Lone pairs nonbonding e - Total e - pairs bonding + lone pairs Lone pairs repel more strongly than bonding pairs!!!
54 A. VSEPR Theory Lone pairs reduce the bond angle between atoms Bond Angle
55 Determining Molecular Shape Draw the Lewis Diagram Tally up e - pairs on central atom (bonds + lone pairs) double/triple bonds = ONE pair Shape is determined by the # of bonding pairs and lone pairs
56 Common Molecular Shapes 2 total Electronic Geometry = linear 2 bond 0 lone BeH 2 LINEAR 180
57 Common Molecular Shapes 3 total 3 bond 0 lone Electronic Geometry = trigonal planar BF 3 TRIGONAL PLANAR 120
58 Molecular Polarity Polar molecule = one end slightly + and one end slightly Molecule with 2 poles = dipolar molecule or dipole
59 Molecular Polarity Shape, symmetry and bond polarity determines molecular polarity H O bond is polar and water is asymmetrical, so H 2 O is polar C Cl bond is polar, but CCl 4 is symmetrical, so molecule is nonpolar
60 Molecular Polarity Identify each molecule as polar or nonpolar O 2 CS 2 CF 4 H 2 O Nonpolar bonds nonpolar Linear nonpolar Tetrahedral nonpolar Bent polar
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