Lecture 06 - Chemical Bonds PDF

What is Chemistry? Lecture 6 CHEMICAL BONDS Semester 1: 2024-2025 Instructor: Nguyen Thi Hoang Hai 1 Open Questions – Physical Properties At room temperature: Why water is a liquid? CO2 is a gas? Table salt is a solid?...

What is Chemistry? Lecture 6 CHEMICAL BONDS Semester 1: 2024-2025 Instructor: Nguyen Thi Hoang Hai 1 Open Questions – Physical Properties At room temperature: Why water is a liquid? CO2 is a gas? Table salt is a solid? Why is table salt dissolved in water? How about iron (nail)? Why is candle wax low melting, soft, and nonconductive, while diamond is high melting and extremely hard? And why is copper shiny, malleable, and able to conduct a current whether molten or solid? 2 http://periodictable.com Open Questions – Chemical Properties Why the iron atoms in the nail can combine with oxygen atoms from O2 in the air to form rust? Can the iron atoms react with N2? Why can we burn propane to get heat? Why not water? 3 Open Questions – Molecular Geometry ❑ Boron trihydride (borane) is planar (flat) ❑ Phosphorus trihydride is pyramidal (pyramid shaped) ❑ General formula AX3 ❑ What makes the difference? ❑ Can we predict it? 4 Open Questions – Answer Chemical bonding (intra-molecular force) Chemical bond ≡ attractive force holding two or more atoms together ➔ Plays a role in determining physical and chemical properties FACT: ~100 chemical elements ➔ millions of chemical compounds formed by chemical bonding 5 Electron configurations (1) is the distribution of electrons into its energy levels and sublevels determines the behavior of the element 6 Electron configurations (2) Period 3 Partial Orbital Diagrams and Electron Configurations 7 Electron configurations (3) Trick to memorizing sublevel filling order: List the sublevels as shown, and read from 1s, following the direction of the arrows. Maximum number of electrons s: 02 e- p: 06 e- d: 10 e- f: 14 e- 8 Electron configurations (4) Quantum numbers of electrons in atoms Name Symbol Permitted Values Property Principal n Positive integers (1, 2, 3, …) Orbital energy (size) Angular l Integer from 0 to n-1 Orbital shape momentum l = 0: s orbital l = 1: p orbital l = 2: d orbital l = 3: f orbital Magnetic ml Integers from –l to 0 to + l Orbital orientation Spin ms +1/2 or -1/2 Direction of e- spin Ex: n = 3, l=2 ➔ 3d orbital Electron orbitals are regions within the atom where electrons have the highest probability of being 9 Electron configurations (5) s subshell p subshell d subshell f subshell ℓ=0 ℓ=1 ℓ=2 ℓ=3 mℓ = 0 mℓ= -1, 0, +1 mℓ= -2, -1, 0, +1, +2 mℓ= -3, -2, -1, 0, +1, +2, +3 One s orbital Three p orbitals Five d orbitals Seven f orbitals 02 s orbital electrons 06 p orbital electrons 10 d orbital electrons 14 f orbital electrons http://chemwiki.ucdavis.edu 10 Octet rule ➔ More stable configuration Octet Rule (for main-group elements) Atoms tend to gain, lose, or share electrons so that each atom has full outermost energy level which is typically 8 electrons (s & p orbitals). First-energy level can contain a max. of 2 electrons 11 Valence electrons (1) ❑ Valence electrons are the electrons in the OUTERMOST energy level (shell) … ➔ most interesting! ❑ B is 1s2 2s2 2p1 ➔ the outer energy level (shell) is 2, and there are 2+1 = 3 electrons in level 2. These are the valence electrons! ❑ Br is [Ar] 4s2 3d10 4p5 ➔ How many valence electrons are present? 12 Valence electrons (2) Number of valence electrons of a main (A) group atom = Group number 13 Covalent chemical bonds H Cl H Cl 14 Lewis structure (1) Lewis electron-dot symbols Two-dimensional structural formula consists of electron-dot symbols that depict each atom and its neighbors, the bonding pairs that hold them together, and the lone pairs that fill each atom's outer level (valence shell). 15 Lewis structure (2) Valence electrons are distributed as ▪ Shared or bonding pairs and ▪ Unshared or lone pairs H ─ Cl Unshared or lone pairs shared or bonding pair This is called a LEWIS structure. 16 Electronegativity Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. ▪ Pauling scale (derived from bond energies) ▪ Pauling Scale: Fluorine: 4.0. ▪ All other elements are assigned values in relation to fluorine 17 Lewis structure (3) Molecular formula Step 1. Place the atoms relative to each other Center position: atom with lowest EN (lower group number, lower EN) Atom placement Step 2. Determine the total number of valence electrons available Sum of valence e- Step 3. Draw a single bond from each surrounding atom to the central atom, and subtract two valence electrons for each bond. Remaining valence e- Step 4. Distribute the remaining electrons in pairs so that each atom ends up with eight electrons (or two for H). Step 5. Cases involving multiple bonds. If, after step 4, a central atom still does not have an octet, make a multiple bond by changing a lone pair from one of the surrounding atoms into a bonding pair to the central atom. Lewis Structure 18 felectron Lewis structure (4) Molecular formula In nearly all their compounds, ❑ Hydrogen atoms form one bond. ❑ Carbon atoms form four bonds. ❑ Nitrogen atoms form three bonds. Atom placement ❑ Oxygen atoms form two bonds. ❑ Halogens form one bond when they are surrounding atoms; fluorine is always a surrounding atom. Sum of valence e- CF2Cl2 CH3OH Remaining valence e- Lewis Structure 19 Lewis structure (5) - Resonance Forms When a single Lewis structure does not adequate reflect the properties of a substance SO2 SO3 NO3- C6H6 (benzene) 20 Lewis structure (6) - Formal Charge Formal charge: is the difference between the number of valence electrons in the free atom and the number assigned to that atom in the Lewis structure Formal charge = (no. of valence e-) – (no. of lone e-) – 0.5*(no. of bonding electron) The formal charges are close to zero as possible Any negative formal charge is located on the most strongly electronegative atom CH4O: CH3OH vs. CH2OH2 21 Lewis structure (7) - Exceptions Electron-deficient molecules (odd number of valence electrons) Ex: NO, NO2, BeF2, BF3 22 Lewis structure (8) - Exceptions Expanded Octets: atoms have d orbitals available for bonding (3d, 4d, 5d) + Distribute the extra electrons around the central atoms as unshared pairs Ex: PCl5, SF6, XeF4 23 Lewis structure (6) Ethylene, C2H4? Table salt, NaCl? 24 Types of bonding Short-range interaction Principles of General Chemistry 25 Types of bonding - ionic Table salt (NaCl) Principles of General Chemistry Atoms (often is metal) transfer electron(s) to other atoms (often non metal), forming oppositely charged ions that attract each other to form a solid. 26 Ionic bonding (1) Formation of ionic bonds in NaCl Bond formed by the transfer of electrons from one atom to another Electrostatic attraction between positive and negative ions 27 Ionic bonding (2) All those ionic compounds were made from ionic bonds. Positive ions and the negative ions are attracted to one another ➔ Ionic compounds are usually between metals and nonmetals (opposite ends of the periodic table). Metal Non-metal 28 Ionic bonding (3) When atoms lose or gain electrons, they acquire a noble gas configuration, but do not become noble gases Na+ and Cl- form an ion pair The net attractive electrostatic forces that hold the cations and anions together are ionic bonds 29 More stable electron configuration for Na Na -- Ne -- - - - - - - - + - - - + - - - - - - - 3s Lose electron 3s Ne Ne Na+: Neon-like configuration ➔ More stable 30 More stable electron configuration for Cl Cl - - Ar -- - - - - - - - - - - - - - - + - - - + - - - - - - - - - - - - - - 3s 3p 3s 3p Cl-: Argon-like configuration ➔ More stable 31 [Na+][Cl-] Both have 8 electrons in their outermost shell This is an ionic compound - - - - - - - - - - - - - - + - - - + - - - - - - - - - - - Na+ Cl- 32 [Na+][Cl-] 33 Ionic bonding (4) The highly ordered solid collection of ions is called an ionic crystal Stoichiometry is an important consideration … 34 Ionic bonding - Summary Electrostatic attraction Oppositely charged ions Metal and a nonmetal Electrons transferred Extremely strong bonds 35 Types of bonding - Covalent Principles of General Chemistry Two atoms share an electron pair localized between their nuclei (shown here as a bond line). Most covalent substances consist of individual molecules, each made from two or more atoms (usually found between nonmetals.) 36 Covalent bonds bond (2) Formed when atoms share electrons ❑ Hydrogen atoms form single bonds H―H ❑ Oxygen atoms form two bonds O=O ❑ Nitrogen atoms form three bonds N≡N ❑ Carbon atoms form four bonds O=C=O 37 Covalent bond (4) Electrons are shared, not transferred Between nonmetals Neutral overall charge Electron orbital expands to include both nuclei Most common type of bond Weaker than ionic bonds 38 Covalent bond (5) - Bond formation A bond can result from an overlap of atomic orbitals on neighboring atoms. H + Cl H Cl Overlap of H (1s) and Cl (3p) Note that each atom has a single, unpaired electron. 39 Covalent bond (6) Covalent bonds exist between atoms that share electrons to form a molecule Double or triple covalent bond This difference is extremely important, and Non-polar covalent bond can indicate how and where a molecule has its Polar covalent bond function. 40 Covalent bond (6) - Number of covalent bonds Multiple covalent bonds ❑more than one pair of e- shared ❑Double bond two pairs of electrons (4 e- total) a double bond is represented by 4 dots or 2 parallel lines ❑Triple bond three pairs of electrons (6 electrons total) a triple bond is represented by 6 dots or 3 parallel lines. 41 Covalent bonds =>  and pi bond Ethylene – C2H4 −  bond −  bond Orbital hybridization 42 Bonding in Ethylene Sigma bond (): electron density between the 2 atoms Pi bond (): e- density above and below plane of nuclei of the bonding atoms 43 Types of covalent bonding Symmetric distribution Asymmetric Distribution 44 Types of bonding - Metallic Principles of General Chemistry Many metal atoms pool their valence electrons to form a delocalized electron "sea" that holds the metal-ion cores together (pure metals) Metal atoms bonded to several other atoms 45 Types of bonding - Metallic 46 Summary on bonding An atom will share, lose or gain enough electrons to become more stable with 8 electrons in its outermost energy level and then it will bond with another atom of a different element. 47 Electronegativity Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. ▪ Pauling scale (derived from bond energies) ▪ Pauling Scale: Fluorine: 4.0. ▪ All other elements are assigned values in relation to fluorine The type of bond can USUALLY be calculated by finding the difference in electronegativity of the two atoms that are going together. Q. ionic, covalent and metallic bond: strongest? weakest? 48 Electronegativity trend Electronegativity values for the elements according to Pauling Trends for electronegativities are the opposite of the trends defining metallic character: Nonmetals have high values of electronegativity, the metalloids have the intermediate values, and the metals have low values. Electronegativity generally increases up a group and across a period. 49 Electronegativity difference ❑ If the difference in electronegativity is between: ▪ 1.7 to 4.0: Ionic ▪ 0.3 to 1.7: Polar Covalent ▪ 0.0 to 0.3: Non-Polar Covalent ❑ NaCl: Na = 0.9, Cl = 3.0 covalent bonding? 50  Molecular Geometry 51 Molecular Geometry VSEPR model Valence Shell Electron Pair Repulsion theory. Most important factor in determining geometry is relative repulsion between electron pairs. Molecule adopts the shape that minimizes the electron pair repulsions. 52 Common geometries Electron pair/group geometry Structural geometry Thẳng Tam giác phẳng Tứ diện Tam giác kim tự tháp đôi Bát diện 53 The VSEPR model The single molecular shape of the linear electron-group arrangement. Beryllium Chloride 0 lone pairs on central atom F Be F 2 atoms bonded to central atom 54 The steps in determining a molecular shape The two molecular shapes of the trigonal planar electron-group arrangement. 55 The steps in determining a molecular shape The three molecular shapes of the tetrahedral electron-group arrangement 56 The steps in determining a molecular shape The four molecular shapes of the trigonal bipyramidal electron-group arrangement phosphorus pentachloride (PCl5) Sulfur tetrafluoride (SF4), 57 The steps in determining a molecular shape The four molecular shapes of the trigonal bipyramidal electron-group arrangement Bromine trifluoride (BrF3), Triiodide ion (I3-) 58 The steps in determining a molecular shape The three molecular shapes of the octahedral electron-group arrangement Sulfur hexafluoride (SF6) iodine pentafluoride (IFs): xenon tetrafluoride (XeF4) 59 The steps in determining a molecular shape Molecular formula Step 1: Write the Lewis structure from the molecular formula to see the relative placement of atoms and the number of electron groups. Lewis Structure Step 2: Count all e- groups around central atom (bonding & nonbonding) Electron-group Arrangement Step 3: Note position of any lone pairs and double bonds Bond Angles Step 4: Draw and name the molecular shape by counting bonding and nonbonding e- groups seperately. Molecular Shape (AXmEn) 60 Structure determination by VSEPR Water, H2O The electron pair geometry is TETRAHEDRAL 2 bond pairs 2 lone The molecular pairs geometry is BENT. 61 Geometry determination Structure of NH3 by VSEPR The electron pair geometry is tetrahedral. lone pair of electrons in tetrahedral position N H H H The MOLECULAR GEOMETRY — the positions of the atoms — is TRIGONAL PYRAMID. 62 Polar and nonpolar molecules A molecule will be polar if: ❑ It has polar bonds. ❑ The center of partial positive charge lies at a different place within the molecule than the center of partial negative charge. Total dipole moment, 𝑑Ԧ 63 Positive Negative Bond polarity HCl is POLAR because it has a positive end and a negative end. (difference in electronegativity) + - Cl has a greater share H Cl in bonding electrons than does H. Cl has slight negative charge (-) and H has slight positive charge (+ ) 64 BOND POLARITY “Like Dissolves Like” ❑Polar dissolves Polar ❑Nonpolar dissolves Nonpolar 65 “Like” Dissolves “Like” Figure 8.34 ❑Ionic salts and polar liquids dissolve better in polar liquids than in nonpolar liquids ❑Nonpolar liquids dissolve better in other nonpolar liquids than in polar liquids 66 Polarity direction of polarity - O - - - O C O H - H carbon dioxide Water (a nonpolar molecule) a polar molecule 67 Polar vs. Nonpolar Covalent Bonds Polar Nonpolar ❑Unequal sharing of electrons ❑Equal sharing of electrons ❑Different elements bonded ❑same elements form a bond different electronegativities same electronegativities ❑Typically shorter bonds ❑Typically longer bonds ❑Stronger bonds (more ionic) ❑Weaker bonds 68 Electron Distributions and Covalent Bonds Symmetric distribution Asymmetric Distribution 69 Electronegativity Polarity is determined by difference in electronegativity – Nonpolar covalent – Polar covalent – Ionic compound 70 Electronegativity and Electronegativity and Chemical chemical Bonds bonds Covalent bond to ionic bond 71 Electronegativity Difference and Bond Type Two identical atoms have the same electronegativity and share a bonding electron pair equally. This is called a nonpolar covalent bond. All homonuclear diatomic molecules have nonpolar covalent bonds: Ex: H2, N2, O2, F2, Cl2, Br2, I2 In covalent bonds between atoms with somewhat larger electronegativity differences, electron pairs are shared unequally. This is called a polar covalent bond Ex: hydrogen chloride gas, HCl. The electrons are drawn closer to the atom of higher electronegativity, Cl With still larger differences in electronegativity, electrons may be completely transferred from metal to nonmetal atoms to form ionic bonds Ex: Sodium chloride, NaCl 72 Electronegativity Difference and Bond Type No sharp cutoff between ionic and covalent bonds. C—H bonds are virtually nonpolar. 73 Ionic vs. Covalent bond 74 75 Comparison of Ionic and covalent compounds 76 Chemical bond ❑ Bond Strength: measured by its dissociation energy Cl2(g) → 2Cl(g); H = +234 kJ HCl(g) → H(g) + Cl(g); H = +431 kJ The HCl bond is stronger than Cl–Cl bond, because it takes more energy to break the bond. ❑ Bond length: assist in determining the overall size and shape of a molecule 77 Chemical bonds ❑Bond lengths → assist in determining the overall size and shape of a molecule ▪ Multiple bond are shorter and stronger than their single bond counterparts ▪ Longer bonds are weaker than shorter bonds ▪ When larger atoms are joined together, the bond becomes longer 78 Lengths of Covalent Bonds Bond Bond Length Type (pm) C-C 154 C=C 133 CC 120 C-N 143 C=N 138 CN 116 Bond Lengths Triple bond < Double Bond < Single Bond 79 Chemical bonds 80 Chemical bonds 81 Hydrogen bond Bond formed when a hydrogen atom is shared between two electronegative atoms 82 Pauling’s Electronegativities Electronegativity Illustrated 83 Lewis XXX structure for O2 + 84 Bonding representation + OR Generally every unpaired electron in the Lewis Dot diagram of an element can form a bond. 85 Ionic bonds (2) Attraction between electrostatic charges is an electrostatic attraction. 86 Ionic bonds (3) Ionic bonding involves 3 steps (3 energies) 1) loss of an electron(s) by one element, 2) gain of electron(s) by a second element, 3) attraction between positive and negative 1) e– 2) Na Cl 3) Na+ Cl– 87