Chemical Bonding & Molecular

Chemical Bonding & Molecular Geometries A word about Electron Configuration The arrangement of an atom’s electrons around its nucleus is called electron configuration. • Quantum number describe the location of electrons in an atom around its nucleus. • Pauli’s Exclusion Principle states that no two electrons can have the same set of quantum numbers. • Electrons will enter the lowest energy level and empty orbital available: Aufbau Principle. Reference credit: pixshark.com Reference credit: pixshark.com Electron Configuration for Nitrogen (N) • • When placing electrons in a orbitals having the same energy (degenerate orbitals), electrons will go singly if possible. – HUND’S RULE The inert gas notation is a shorthand for writing the electron configuration. – Find the noble gas preceding element (full shell); this is referred to as the CORE electrons (e.g. [He]) – The VALENCE electrons (e.g. 2s2 2p3 are the electrons outside of the core electrons. • These are the electrons that participate in bonding and forming ions • • Atomic number = 7 Electron configuration for N – 1s2 2s2 2p3 = total electrons – Shorthand: [He] 2s2 2p3 Box orbital diagram of N 2p E 2s 1s Periodic Table & Electron Configuration Reference credit: maddchem.weebly.com Role of Valence in Bonding • The outermost electrons surrounding the atom’s nucleus are referred to as the valence electrons. – Valence electrons are in the outermost energy shell & subshells (i.e., highest energy level). • It is the valence electrons that are participants in bonding, forming ions, etc. • All electrons that are not valence electrons are known as the core electrons. – Core electrons do not participate in chemical reactions, bonding, formation of ions. Problem: Given below are the electron configurations for the elements O, S, Se & Te: O: 1s22s22p4 S: 1s22s22p63s23p4 Se: 1s22s22p63s23p64s23d104p4 Te: 1s22s22p63s23p64s23d104p65s24d105p4 1. Write the inert gas shorthand electron configuration & identify the valence electrons. 2. What feature makes them have similar chemical properties? Electron Configuration for Cations Alkali & Alkaline Earth (Group 1(A) & 2(A)): • s column metals when forming cations will lose their s-valence electrons. • Obtain the “noble look” • Example: • Ba: [Xe] 5s2 • Ba2+ : [Xe] Main Group or “p” block elements (3A-7A or 1317): • Will lose their p-valence electrons FIRST and then their s-valence electrons. • Examples: • Pb: [Xe] 4f14 5d10 6s2 6p2 Pb2+: [Xe] 4f14 5d10 6s2 • Pb: [Xe] 4f14 5d10 6s2 6p2 Pb4+: [Xe] 4f14 5d10 • Full d & f subshells are very stable. – They behavior as “pseudo inner (full shells) core electrons. Transition metals, the d” block elements (1B-2B or 3-12): • • When forming cations, the s-valence electrons are lost FIRST and then the dvalence electrons. Examples: • Fe [Ar]3d6 4s2 • Fe [Ar]3d6 4s2 Fe2+ [Ar]3d6 Fe3+ [Ar]3d5 • Inner transition metals, the “f” group elements (lanthanides #58-71 & actinides #90 -103): • The s-valence electrons are lost FIRST and then the f-valence electrons. • Examples: • Am: [Rn]5f77s2 • Am2+ : [Rn]5f7 • Am6+: [Rn]5f3 • Half-filled d and f subshells are very stable. Ion Electron Configuration: Anion • Nonmetal elements tend to gain electrons to obtain the “noble look.” – Form anions by gaining (adding) electrons in their valence shell. • Examples: P: [Ne] 3s2 3p3 P3-: [Ne] 3s2 3p6 or [Ar] I: [Kr] 4d105s2 5p5 I-: [Kr] 4d105s2 5p6 or [Xe] Problems: 1. Identify the element having the electron configuration of (1s22s22p63s23p2). 2. Write the electron configuration for Ni(II) ion. 3. Write the electron configuration for Se2-. 4. Write the electron configuration for C4+ . 5. Write the electron configuration for Mo. 6. Write the electron configuration and identify the valence electrons for: a. Mn2+ b. O2c. Sr Why Do Atoms Bond? • Atoms loose or gain electrons to achieve a full shell (i.e., noble) look; low energy state. – Bonding is another way atoms can achieve the full shell look. • Bonding occurs when atoms either exchange their valence electrons (e.g., ionic) or share their valence electrons (e.g., covalent and metallic) electrons between atoms. – The formed bond maximizes the attraction forces & minimizes the repulsive forces. • Categories of bonding: – Ionic – Covalent – Metallic • The strength of the bond between the atoms depends on many factors such as type of bond formed, element electronegativity, size of atoms and the number of bonds between atoms. Ionic Bonding: Metal & Nonmetal Interaction • Ionic Bonding: – Is the complete transfer of one or more electrons from one atom (metal) to another atom (nonmetal). • The resulting attraction force in the bond is between two charges (i.e., cation + anion). – Ionic bond strength can be calculated using Coulomb’s Law ( F= k 𝑞1 𝑞2 ). 𝑟2 Reference credit: imgacarade.com Electronegativity • Is a chemical property of an atom’s ability to attract electrons (electron density towards itself when bonded to another element. • Trend-wise: Electronegativity: INCREASES ACROSS a period and DECREASES DOWN a column. • Electronegativity numerical values range from 0.07 (metals) to 4.0 (nonmetals). • The attraction or pulling on electrons within the bond causes a separation of charge within the bond; referred to as a DIPOLE MOMENT (d). • The greater the difference in attraction between the elements in the bond the more POLAR the bond . • As the electronegativity difference increases between two atoms, the bond becomes more ionic. Reference credit: OpenStax2017 • The distribution of electron density in the HCl molecule is uneven. • The electron density is greater around the chlorine nucleus. Electronegativity & Bond Types: Bonding Continuum • As the electronegativity difference increases between two atoms, the bond becomes more ionic. • The magnitude of ∆EN predicts the amount of charge separation for the bonds that share electrons unevenly. • There are no sharp cutoffs between nonpolar, polar, and ionic bonds. • There is a continuum of bonding between 100% even sharing of electrons and 100% transfer of electrons. Reference credit: Interactive General Chemistry Electronegativity Values Reference Credit: OpenStax2017 • • The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table. The most electronegative element is fluorine (4.0) Covalent Bonding: Nonmetal to Nonmetal Attraction • Covalent bonding occurs when two nonmetal atoms bond together by sharing valence electrons between themselves. – The bond arises from the mutual attraction of two nuclei for the same electrons. • Bond is a balance between the attractive forces (i.e., proton (+) to electron (-) and two types of repulsive forces (i.e. Electron (-) to electron (-) and (proton (+) to proton (+)). • Bonds formed are result of overlapping atomic orbitals on neighboring atoms. Reference credit: imgarcarade Metallic Bonding: Metal with Metal • Metallic bonding is electron sharing AMONG metallic atoms. – Metals give up their valence electrons (i.e., forming metal cations) to share among the other metal ions. • The bond is a balance between the attractive forces (i.e., proton (+) to electron (-) and two types of repulsive forces (i.e. Electron (-) to electron (-) and (proton (+) to proton (+)). • Metallic bonding electron sharing differs from typical covalent bonding because it is communal sharing of electrons (e.g., sea of electrons) among the metal cations and NOT a one on one between metal atoms. – Explains the properties of metals & their alloys. Reference credit: chemistry.tutorvista.com Lewis Symbols and Structures: An Introduction • Lewis structures are the “electron configurations” of atoms and or ions. – They illustrate how the atoms’ valence electrons forms connections. – Lewis structure are drawn mostly for covalent molecules, polyatomic ions and alkali and alkaline earth ionic compounds. • Lewis symbol is drawn by using the element’s symbol surrounded by its valence electrons. Reference credit: OpenStax2017 Reference credit: chemwiki.ucdavis.edu Lewis Dot Structures of Ionic Compounds • There are several pictorial representations of Lewis Dot Structures of ionic compounds. – Ionic compound bonding is the result of electron transfer between atoms; specifically, a metal and a nonmetal (i.e., cation-anion). Reference: www2maxwellsyr.edu Reference: www.ilectureonline.com Lewis Structures of Molecules • Lewis Theory states: – When atoms bond, they tend to gain, lose, or share electrons to result in a “full shell” or noble gas configuration. • Lewis structures of molecule and ions illustrates the distribution of VALENCE ELECTRONS. • The valence electrons are distributed as either SHARED PAIRS (BONDED) or UNSHARED( LONE) PAIRS within the molecule. lone pair • • H • • Br • • shared or bonded pair H • • • • lone pair • • Br • • shared or bonded pair An Overview How to Draw a Lewis Structure • • When drawing the Lewis structure for a simple molecule or ion the focus is on the center atom. – The center atom is the atom to which all other atoms will be bonded. • In a molecular formula the center atom is usually the first nonmetal written in the formula. Draw a skeletal structure. – Attaching all atoms to the center atom. • Count the number of valence electrons each atom is bring into bond. • Arrange the valence electrons satisfying the criteria as given by the Octet Rules & Exceptions to the Octet Rules. – Example: • Hydrogen atoms can not be center atoms for they can only have two electrons as dictated by quantum theory. • Second period nonmetals must have eight electrons surrounding the atom whether they are a center atom or a attaching atom. • Third period nonmetals must have eight electrons around the atom as an attachment BUT can have more than eight electrons if they are a center atom. An Overview How to Draw a Lewis Structure • Multiple Bonding Elements: – Double Bond (represented by two bond pairs or two lines between atoms). • A double bond consists of 4 electrons or 2 bonding pair of electrons between two atoms. • C, N, O, S & P can double bond with each other and themselves. – Triple Bonds (represented by three bond pairs or three lines between atoms). • A triple bond consists of 6 electrons or 3 bonding pair of electrons between two atoms. • C, N, O & S can triple bond with each other and themselves. Example of Drawing a Lewis Structure Draw the structure for water, H2O. 1. O has to be the center atom because of its limitation of only forming a single bond. 2. Draw the skeletal structure. O H H 3. Count the number of valence electrons each atom is bring into the compound. Atom Valence Electrons 1 Oxygen 6 electrons 6 electrons 2 Hydrogen 1 valence each or 2 x 1 = 2 electrons TOTAL: 8 valence electrons 4. Arrange valence electrons within the structure according to criteria. Water has an electron pair geometry that is tetrahedral & molecular geometry that is called bent. O H H Examples: Exceptions to the “Octet (Deficient & Hypervalent) Problem: Draw the Lewis dot structure for: 1. NH3 2. BF3 3. CH2O 4. HF 5. HCN 6. I37. ICl4- 8. BrCl3 Resonance Structures & Formal Charges • Resonance structures: when more than one valid Lewis structure can be drawn for a molecule. – Often seen when have elements in compound that can multiple bond. • Formal charges is a method to predict which resonance structure is the more stable (i.e., energy favored) version. – Formal charges are hypothetical charges on an atom in a molecule or ion. • It is a bookkeeping procedure. Resonance Structures Example: NO2-, nitrite ion. • Draw all possible Lewis structures that met the criteria. – Total number of valence electrons: 18 • N: 5 electron • O: 6 x 2 12 electron • Anion: 1 electron – Nitrogen is center atom • Nitrogen & oxygen atoms must have 8 electrons around them. • Two legitimate Lewis dot structures can be drawn for the nitrite ion. – RESONANCE STRUCTURES. Reference credit: chemistry.about.com Geometries of Molecules & Ions: VSEPR Valence Shell Electron Pair Repulsion theory. • Predicts the molecular structure of molecular compounds and polyatomic ions. – Assume compound’s valence electrons around the CENTER atom reflects an arrangement that favors minimal electron repulsion. • Most important factor in determining geometry is relative repulsion between electron pairs. • Must be able to draw Lewis structures of compounds. – Theory does not work for uneven valance electron molecules. • Example: NO Electron-Pair Geometry (family) vs. Molecular Geometry (shape) Molecular Geometry (shape) Electron-Pair Geometry (family) • The geometry that includes all the electron pairs. – Classification of the geometric structure (i.e. family). • Family Groupings: – Linear – Trigonal planar – Tetrahedral/tetrahedron – Trigonal bipyramidal – Octahedral • The geometric structure includes only the placement of atoms in compound. – Gives the molecule or ion its shape. – Takes into account the influence of lone pairs around the center atom on molecular structure. • Angles – Predict molecular polarity. FAMILY Linear Trigonal planar MOLECULAR SHAPE Electron pair geometry (family) & Molecular geometry (shape) Tetrahedral Trigonal bipyramidal Octahedral Reference credit: Socarates.org Linear Electron-pair geometry (family) (2) • Molecular Geometry (shape) – Linear • Bond angle: 180 o Reference credit: people.uwplatt.edu • Usually 2 atoms in compound can be three, but will have a linear geometry. Reference credit:OpenStax2017 Trigonal Planar Electron-pair geometry (family) (3) • Molecular Geometry (shape) – Trigonal planar • No lone pairs, all attachments. • Bond angle: 120o – Bent • One lone pair & two attachments. • Bond angle: <120o Reference credit: uwplatt.edu Reference credit: www.learner.org Tetrahedral Electron-pair geometry (family) (4) • Molecular Geometry (shape) – Tetrahedral • No lone pairs, all attachments. • Bond angle: 109 o – Trigonal pyramidal • One lone pair & three attachments • Bond angle: <109 o – Bent, Angular, V-shape • Two lone pairs & two attachments • Bond angle: <109 o Reference credit: quizlet.com Trigonal Bipyramidal Electron-pair geometry (family) (5) – Trigonal bipyramidal • No lone pairs, all attachments. • Bond angle: 90o & 120o – Sawhorse or Seesaw • One lone pair & four attachments. – Lone pair will occupy the 120 o area – T-shape • Two lone pairs & three attachments. – Lone pair will occupy the 120 o area – Linear • Three lone pairs & two attachments. – Lone pair will occupy the 120 o area •• • Molecular Geometry (shape) Reference credit: quizlet.com Octahedral Electron-pair geometry (family) (6) • Molecular Geometry (shape) – Octahedral • No lone pairs all attachments. • Bond angle: 90o – Square pyramidal • One lone pair & five attachments. • Bond angle: >90o – Square planar • Two lone pairs & four attachments. – Bond angle: >90 o – T-shape • Three lone pairs & three attachments. – Bond angle: >90 o – Linear • Four lone pairs & two attachments. – Bond angle: >90 o Reference credit: www.cram.com Problem: – Draw the Lewis structures – Identify the electron-pair (family) geometries – Determine the molecular (shape) geometries 1. 2. 3. 4. 5. 6. NH4+ XeF4 OCS SbH3 PCl3 KrF2 Lone Pair Effect on Molecular Geometries Reference credit: www.mhhe.com • Lone pair electrons require more space around the center atom. – Their electron density is exclusively on the central atom rather than shared like bonding electron groups. – This affects the bond angles, making them smaller than expected. • Electron-par geometries (e.g., linear, trigonal planar, tetrahedral and octahedral) where all bond angles are equal DO NOT have positions around the center atom that can minimize the electron  → electron repulsion. • Molecules with lone pairs or different kinds of surrounding atoms will have distorted bond angles and different bond lengths, but the shape will be a derivative of one of the basic shapes. Molecular Polarity: Molecular Geometries Criteria for a polar MOLECULE: • 1. Polar covalent bonds MUST be present in the molecule. • Electronegativity difference. – Dipole moments measured. 2. Asymmetric distribution of electron density throughout the molecule. • Molecular geometries: – Unsymmetrical shape. – Nonbonding electron pairs affect molecular polarity; they pull the electron density strongly. The polarity affects the intermolecular forces of attractions, which in turn influences the physical properties (e.g.. boiling point, solubility, melting point) of the compound. Reference credit: wps prenhall.com CO2 vs. H2O: Polar or Not? • Both molecules contain polar covalent bonds. – Met one of the criteria for determining if a molecule is polar. Reference credit: Openstax2017 • • CO2 is nonpolar – Sum of dipole moments is zero. • Geometry give a symmetric distribution of the electron density throughout the molecule. H2O is polar – Sum of the dipole moments is not zero. – You have an asymmetric distribution of electron density throughout molecule due to the geometry of the molecule and lone pairs on the center atom. Problem: Determine the polarity of these molecules. a. sulfur hexafluoride, SF6 b. bromine pentafluoride, BrF5 c. Ammonia, NH3 Overview of Bond Properties • Bond Properties – Length: • Distance between to atom nuclei in a bond. – Strength: • For covalent molecules and ions, strength is related to the energy needed to break the bond (i.e., the energy required to separate the atoms in the bond. • For ionic compounds, strength is related to the attraction forces between the cation and anion or in other words, the lattice energy. – Energy: • Energy associated with the breaking of bonds, or the energy required to separate atom or ions in a bond. Factors that Affect the Length of a Bond Bond length is the distance between two atom nuclei. Reference credit: chemwiki.ucdavis.edu Reference credit: ch301.cm.utexas.edu Factors that influence bond length: • Size of the atoms in the bond. – In ionic compounds the size of the cation & anion. – In molecular compounds the size of the nonmetal atom nuclei • Length is measured in Angstroms – 1 A = 1 x 10-2 pm Reference credit: chemwiki.ucdavis.edu Factors that Affect the Length of a Bond • Ionic charge: – In ionic compounds the ionic charge along with ion size affects the distance. • Electronegativity: – In molecular compounds and ions, electronegativity of the atoms in the bond. Reference credit: www.keyword-suggestions.com Factors that Affect the Length of a Bond • Number of electrons between the atoms. • Multiple bonds – Single > double > triple Reference credit: www.boundlesss.com Factors that Affect Bond Strength: Covalent Bond Strength: – For covalent molecules and ions, strength is related to the energy needed to break the bond (i.e., the energy required to separate the atoms in the bond. • Greater the number of bonds between atom nuclei the stronger the bond. • Example: – C≡C (837 kJ) > C=C (611 kJ) > C−C (347 kJ) – C≡N (891 kJ) > C=N (615 kJ) > C−N (305 kJ) • The shorter the bond length the stronger the bond. • Example: – Br−F (237 kJ) > Br−Cl (218 kJ) > Br−Br (193 kJ) Factors that Affect Bond Strength: Ionic Bond Strength: • For ionic compounds, strength is related to the attraction forces between the cation and anion or in other words, the lattice energy. • Lattice energy: Heat + MX (s) → Mn+ (g) + Xn- (g) – Energy required to break the bond. • Ionic bonds are the strongest bonds. – Electrostatic attraction. • Coulombic Forces – Closer the two charges the stronger the force. 𝑞1 𝑥 𝑞2 – F=k 𝑑2 Reference credit: www.imgarcade.com