Chemistry Lecture: Chemical Bonds - PDF

Chemistry Lecture Basic Medical Course Chemical Bonds László Virág Department of Medical Chemistry Faculty of Medicine University of Debrecen www.medchem.unideb.hu email: [email protected] ...

Chemistry Lecture Basic Medical Course Chemical Bonds László Virág Department of Medical Chemistry Faculty of Medicine University of Debrecen www.medchem.unideb.hu email: [email protected]  Chemical bonds We consider three main types of chemical bonds: ionic bond (electrostatic forces which hold ions together, e.g. NaCl); covalent bond (results from sharing electrons between atoms, e.g. Cl2); metallic bonding (refers to metal nuclei floating in a sea of electrons, e.g. Na). In all chemical bonds, electrons are shared or transferred between atoms.  Core electrons not involved in bonding  VALENCE ELECTRONS: The electrons in the outermost shell are called VALENCE ELECTRONS. For the main group elements, the group number is the number of valence electrons. How many valence electrons in Ca? N? O? What is their electron configuration? ©  OCTET RULE: What do we know about the Noble Gases? They are stable or inert. That is, they don’t react. Their valence shell is full. Compounds form because elements „want to achieve” a full valence shell (or the electron configuration of a noble gas). This is the OCTET RULE. Elements „want to achieve” eight valence electrons (or two for helium). Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons ✓C would like to gain 4 electrons ✓N would like to gain 3 electrons ✓O would like to gain 2 electrons  OCTET RULE: How can elements achieve an octet? Bonding TRANSFER electrons to form IONIC bond. SHARE electrons to form COVALENT bond.  IONS When elements lose or gain electrons, they form IONS. Ions have a positive charge (loss of e-) or a negative charge (gain of e-) equal to the number of electrons lost or gained.  POSITIVE IONS: Metals (left side of periodic table) will lose electrons to form positive ions. Positive ions are called CATIONS. A (monatomic) metal ion (cation) is named by its element name.  NEGATIVE IONS: Nonmetals (right side of periodic table) will gain electrons to form negative ions. Negative ions are called ANIONS. A (monatomic) anion is named by placing -ide at the end of the root of the element’s name.  Ionic Bond Combining the above processes, the overall effect is the transfer of ONE electron from sodium to chlorine.  Describing Ionic Bonds The atom that loses the electron becomes a cation (positive). + Na([Ne]3s ) → Na ([Ne]) + e 1 - The atom that gains the electron becomes an anion (negative). − Cl([Ne]3s 3p ) + e → Cl ([Ne]3s 3p ) 2 5 - 2 6  Ionic Bond An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions. This type of bond involves the transfer of electrons from one atom (usually a metal) to another (usually a nonmetal). The number of electrons lost or gained by an atom is determined by its need to be “isoelectronic” with a noble gas.  Describing Ionic Bonds Consider the transfer of valence electrons from a sodium atom to a chlorine atom. + − Na + Cl → Na + Cl e- The resulting ions are electrostatically attracted to one another. The attraction of these oppositely charged ions for one another is the ionic bond. Ionic Bonds: One Big Greedy Thief Dog! 1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.  Ionic Bonding Energetics of Ionic Bond Formation The formation of Na+(g) and Cl-(g) from Na(g) and Cl(g) is endothermic. Why is the formation of NaCl(s) exothermic? The reaction NaCl(s) → Na+(g) + Cl-(g) is endothermic (H = +788 kJ/mol).  Ionic Bonding Energetics of Ionic Bond Formation The energy required to separate one mole of ions in an ionic lattice into gaseous ions is called the lattice energy, Hlattice. Lattice energy depends on the charge on the ions, and the size of the ions. The specific relationship is given by Coulomb’s equation: QQ E=k 1 2 d as Q1 and Q2 increase, E increases; as d increases, E decreases.  Na+ and Cl- Ions form a 3D Lattice Lattice energy cannot be determined experimentally due to the difficulty in isolating gaseous ions. The energy value can be estimated using the Born-Haber cycle, or it can be calculated theoretically with an electrostatic examination of the crystal structure.  Ionic Bonding Energetics of Ionic Bond Formation The Born-Haber cycle  Ionic Bonding Born-Haber Cycle This is a thermodynamic cycle that analyzes lattice energy precisely. The Born-Haber cycle looks at the formation of NaCl(s) from Na(s) and Cl2(g). The direct route is Hf: Na(s) and ½Cl2(g) → NaCl(s), H = -410.9 kJ Alternatively, we can form sodium gas (endothermic), then chlorine atoms (endothermic), then sodium ions (ionization energy for Na, endothermic), then chloride ions (electron affinity for Cl, exothermic), then form the ionic lattice (exothermic).  Ionic Bonding Energetics of Ionic Bond Formation  Ionic Bonding complete transfer of electrons between atoms converts them to ions, and they will form ionic compounds Ionic compounds tend to form between metals and non-metals Ionic compounds form extensive arrays called crystals Sodium and chloride are held together in the crystal by ionic bonds (strong electrostatic interactions) Covalent Bonding  Covalent Bonding In ionic bonding one atom completely loses an electron while the other gains the electron. When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. When similar atoms bond, they share pairs of electrons to each obtain an octet. Each pair of shared electrons constitutes one chemical bond. Example: H + H → H2 has electrons on a line connecting the two H nuclei.  Covalent Bonds When two nonmetals bond, they often share electrons since they have similar attractions for them. This sharing of valence electrons is called the covalent bond. These atoms will share sufficient numbers of electrons in order to achieve a noble gas electron configuration (that is, eight valence electrons). The tendency of atoms in a molecule to have eight electrons in their outer shell (two for hydrogen) is called the octet rule.  Covalent Bond Between nonmetallic elements of similar electronegativity. Formed by sharing electron pairs Stable non-ionizing particles, they are not conductors at any state Examples; O2, CO2, C2H6, H2O, SiC Covalent Bonding Energy of interaction of two hydrogen atoms  Covalent Bonding Multiple Bonds It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): One shared pair of electrons = single bond (e.g. H2); Two shared pairs of electrons = double bond (e.g. O2); Three shared pairs of electrons = triple bond (e.g. N2). H H O O N N Generally, bond distances decrease as we move from single through double to triple bonds.  NONPOLAR COVALENT BONDS when electrons are shared equally H2 or Cl2 Non-Polar Covalent Bond  POLAR COVALENT BONDS when electrons are shared but shared unequally H2O - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen. Bond Polarity and Electronegativity In a covalent bond, electrons are shared. Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. There are some covalent bonds in which the electrons are located closer to one atom than the other. Unequal sharing of electrons results in polar bonds. Polar Covalent Bonds: Unevenly matched, but willing to share.  Polar Covalent Bonds Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, electronegativity increases from the lower-left corner to the upper-right corner of the periodic table. The current electronegativity scale, developed by Linus Pauling, assigns a value of 4.0 to fluorine and a value of 0.7 to cesium. Bond Polarity and Electronegativity Electronegativity  Electronegativities of the Elements 9_12 H 2.1 IA IIA IIIA IVA VA VIA VIIA Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl VIIIB 0.9 1.2 IIIB IVB VB VIB VIIB IB IIB 1.5 1.8 2.1 2.5 3.0 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br 0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I 0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5 Cs Ba La–Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At 0.7 0.9 1.1–1.2 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2 Fr Ra Ac–No 0.7 0.9 1.1–1.7 Bond Polarity and Electronegativity Electronegativity and Bond Polarity Difference in electronegativity is a gauge of bond polarity. There is no sharp distinction between bonding types. The positive end (or pole) in a polar bond is represented + and the negative pole -.  Polar Covalent Bonds The absolute value of the difference in electronegativity of two bonded atoms gives a rough measure of the polarity of the bond. When this difference is small (less than 0.5), the bond is nonpolar. When this difference is large (greater than 0.5), the bond is considered polar. If the difference exceeds approximately 1.8, sharing of electrons is no longer possible and the bond becomes ionic. Electronegativity and bond type Just as a summary to what each bond looks like… Lewis dot structures  Electron Electron distribution Distribution is depicted with Lewis (electron dot) in Molecules structures This is how you decide how many atoms will bond covalently! (In ionic bonds, it was G. N. Lewis 1875 - 1946 decided with charges) Bond and Lone Pairs Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. H Cl lone pair (LP) shared or bond pair This is called a LEWIS structure.  Lewis Electron-Dot Symbols A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element. Group I. Group II.. Group III Group IV Group V Group VI Group VII Group VIII Na.. Mg..Al.. Si.. : : : : : :.. P :S.. : Cl. : Ar: Note that the group number indicates the number of valence electrons.  The Lewis Dot Structure Lewis devised a system for determining chemical bonding that involves representing each valence electron as a dot. For Ionic Compounds... the metal will lose all of it’s ‘dots’ so that it is isoelectronic with the noble gas in the previous row. The non-metal will gain electrons until it has eight electrons. For Covalent Compounds... The electrons placed in- between the atoms are shared and counted by both atoms. Both atoms must have eight electrons. Lewis Electron-Dot Formulas (ionic compound) : : : F.. Mg. : :. F: - 2+ - : : : : [: F: ] Mg [: F: ]  Lewis Structures You can represent the formation of the covalent bond in H2 as follows: H. +.H : H H  Lewis Structures The shared electrons in H2 spend part of the time in the region around each atom. H H: In this sense, each atom in H2 has a helium configuration.  LEWIS STRUCTURES In writing Lewis structures, only the valance electrons are used. There are two kinds of electron pairs: – Shared electrons form covalent bonds (indicated by a line) – Unshared pairs of electrons (indicated by two dots)  Lewis Structures Covalent bonding in a molecule is repre- sented by a Lewis structure. A valid Lewis structure should have an octet for each atom except hydrogen. H2 H + H → H H or H H → Bonding Cl + Cl2 Cl Cl Cl electrons or Cl Cl Nonbonding electrons  Drawing Lewis Structures Sum the valence electrons from all atoms. Add one for each negative charge and subtract one for each positive charge. Identify the central atom (usually the one with the highest atomic mass and closest to the center of the periodic table). Place the central atom in the center of the molecule and add all other atoms around it. Place one bond (two electrons) between each pair of atoms. Complete the octets for atoms bound to the central atom. Place leftover electrons on the central atom. If there are not enough electrons to give the central atom an octet, try multiple bonds.  Lewis Structures Draw Lewis structures for: HF H F or H F H O H H 2O or H O H NH3 H N H or H N H H H H H CH4 H C H or H C H H H  Double and Triple Bonds Atoms can share four electrons to form a double bond or six electrons to form a triple bond. O2 O =O N2 N N The number of electron pairs is the bond order.  Drawing Lewis Structures O COCl2 24 ve’s Cl C Cl HOCl 14 ve’s H O Cl − O ClO3− 26 ve’s O Cl O H CH3OH 14 ve’s H C O H H  Delocalized Bonding: Resonance The structure of ozone, O3, can be represented by two different Lewis electron-dot formulas. : : O or O : : : : : : : : O O: :O O Experiments show, however, that both bonds are identical.  Exceptions to the Octet Rule Although many molecules obey the octet rule, there are exceptions where the central atom has more than eight electrons. Generally, if a nonmetal is in the third period or greater, it can accommodate as many as twelve electrons, if it is the central atom.  Exceptions to the Octet Rule There are three classes of exceptions to the octet rule: Molecules with an odd number of electrons; Molecules in which one atom has less than an octet; Molecules in which one atom has more than an octet. Odd Number of Electrons Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons. N O N O  Exceptions to the Octet Rule Less than an Octet Relatively rare. Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. Most typical example is BF3. Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.  Exceptions to the Octet Rule More than an Octet This is the largest class of exceptions. Atoms from the 3rd period onwards can accommodate more than an octet. Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.  FORMAL CHARGE Often it is possible to write two different Lewis structures for a molecule that: – differ in the arrangement of their atoms – and both satisfy the octet rule However, only one of the two structures actually exists in nature. So, you must have a way of determining which of the two structures is more plausible.  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. Mathematically: Cf = X – (Y + Z/2) Where X = the number of valance electrons in the free atom (Group Number) Y = the number of unshared electrons Z = the number of shared electrons  FORMAL CHARGE The more likely Lewis structure is the one that: – has formal charges as close to zero as possible and – has negative formal charge located on the most strongly electronegative atom.  METALLIC BOND bond found in metals holds metal atoms together very strongly  Metallic bonding Metal cations are immersed in a sea of delocalized electrons. Electron cloud shields positively charged ion cores from mutually repulsive electrostatic forces Nondirectional bond Weak or strong Good conductors for electricity and heat (free electrons). Examples; Na, Fe, Al, Au, Co Metallic Bonds: Mellow dogs with plenty of bones to go around. Metallic Bond, A Sea of Electrons  Metals Form Alloys Metals do not combine with metals. They form Alloys which is a solution of a metal in a metal. Examples are steel, brass, bronze and pewter. Fe+C; Cu+Zn; Cu+Sn(+Al, Mn, Ni, Zn); Sn+Sb(+Cu+Bi)