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Topic 4 Bonding

Topic 4 Bonding. 4.1 Ionic Bonding Valence Electrons. Valence electrons are the electrons in the highest occupied energy level of an element’s atoms. The number of valence electrons largely determines the chemical properties of an element. Determining the Number of Valence Electrons.

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Topic 4 Bonding

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  1. Topic 4 Bonding

  2. 4.1 Ionic BondingValence Electrons • Valence electrons are the electrons in the highest occupied energy level of an element’s atoms. • The number of valence electrons largely determines the chemical properties of an element.

  3. Determining the Number of Valence Electrons • To find the number of valence electrons in an atom of a representative element, simply look at its group number. • Atoms of the Group 1A elements (hydrogen, lithium, sodium, and so forth) all have one valence electron, corresponding to the 1 in 1A. • Carbon and silicon atoms, in Group 4A, have four valence electrons. • The noble gases (Group 0) are the only exceptions to the group-number rule: Atoms of helium have two valence electrons, and atoms of all the other noble gases have eight valence electrons.

  4. Valence electrons are usually the only electrons involved in chemical bonds. • As a general rule, only the valence electrons are shown in electron dot structures. • Electron dot structures are diagrams that show valence electrons in the atoms of an element as dots.

  5. The Octet Rule • The octet rule states that in forming compounds, atoms tend to achieve the electron configuration of a noble gas. • An octet is a set of eight. • Atoms of each of the noble gases (except helium) have eight electrons in their highest occupied energy levels • Atoms of metals tend to lose their valence electrons, leaving a complete octet in the next-lowest energy level. • Atoms of nonmetals tend to gain electrons

  6. Ionic Compounds • A cation (+) is formed when an atom loses electrons. Usually metals are cations. • An anion (-) is formed when an atom gains electrons. Usually nonmetals are anions • Cations and anions have opposite charges and are attracted to one another. • These attractive forces can hold the ions together in an ionic bond, forming a compound. • Ionic compounds are usually made up of a metal and nonmetal.

  7. Formulas of ionic compounds • A chemical formula represents which & how many atoms are in a compound. • Cations and anions bond so that the compound has no net charge overall. The positive and negative charges cancel each other out. The positive ion is listed first in the formula. Example: Fe2+ Cl1- FeCl2 the 2 is an example of a subscript

  8. Naming Ionic Compounds • Name the cation 1st & the anion 2nd. • Monatomic (1 element) cations use the element name. • Monatomic anions use the root of their element name plus the suffix ide • If the compound has a polyatomic ion, simply name that ion without any change. • If the ion has more than one “common” form, it has to be labeled with a Roman numeral. Ex. Iron (II) phosphate

  9. Polyatomic ions to know(p. 69 IB book) NO31- nitrate OH1- hydroxide HCO31- hydrogen carbonate SO42- sulfate CO32- carbonate PO43- phosphate NH41+ ammonium

  10. Ionic compounds have a lattice structure • Most ionic compounds are crystalline solids at room temperature. • The component ions in such crystals are arranged in repeating three-dimensional patterns.

  11. Molecular (covalent) compounds • A covalent bond results from the sharing of electrons. The octet rule still applies • Covalent bonds generally occur when elements are close to each other on the periodic table. • The majority of covalent bonds form between nonmetallic elements. (remember that ionic bonds form between metals and non metals)

  12. Diatomic Elements(elements that exist in pairs) • Hydrogen H2 • Oxygen O2 • Nitrogen N2 • Fluorine F2 • Chlorine Cl2 • Bromine Br2 • Iodine I2

  13. Naming Molecular (covalent) Compounds • The 1st element is named first, using the entire element name • The 2nd element is named using the root of the element & adding the suffix –ide • Prefixes are used to indicate the # of atoms of each type that are present. *Exception: The 1st element in a formula never uses the prefix mono-

  14. Prefixes in Molecular (covalent) Compounds

  15. Examples NH3 nitrogen trihydride N2H4 dinitrogen tetrahydride H2O dihydrogen monoxide (common name water)

  16. Single Covalent Bonds • Two atoms held together by sharing one pair of electrons are joined by a single covalent bond. • An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots. • The pair of shared electrons forming the covalent bond is also often represented as a dash, as in H—H for hydrogen.

  17. Double & Triple Covalent Bonds • A double covalent bond is a bond that involves two shared pairs of electrons. • Similarly, a bond formed by sharing three pairs of electrons is a triple covalent bond.

  18. Bond Polarity • Covalent bonds differ in terms of how the bonded atoms share the electrons. • The bonding pairs of electrons in covalent bonds are pulled between the nuclei of the atoms sharing the electrons.

  19. Nonpolar covalent bond • When the atoms in the bond pull equally (as occurs when identical atoms are bonded), the bonding electrons are shared equally, and each bond formed is a nonpolar covalent bond.

  20. A polar covalent bond, known also as a polar bond, is a covalent bond between atoms in which the electrons are shared unequally. • The more electronegative atom attracts more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. • Reminder that electronegativity increases from left to right and from bottom to top of the periodic table.

  21. Example of polar bond • Hydrogen has an electronegativity of 2.1, and chlorine has an electronegativity of 3.0. • The chlorine atom, with its higher electronegativity, acquires a slightly negative charge. • The hydrogen atom acquires a slightly positive charge. δ+ δ– H—Cl The lowercase Greek letter delta (δ) denotes that atoms in the covalent bond acquire only partial charges, less than 1+ or 1–.

  22. These partial charges are shown as clouds of electron density. • This electron-cloud picture of hydrogen chloride shows that the chlorine atom attracts the electron cloud more than the hydrogen atom does. .

  23. G. N. Lewis 1875 - 1946 Electron Distribution in Molecules • Electron distribution is depicted withLewis electron dot structures • Valence electrons are distributed as shared orBOND PAIRS (BP) and unshared orLONE PAIRS (LP).

  24. •• H Cl • • •• lone pair (LP) shared or bond pair Bond and Lone Pairs • Valence electrons are distributed as shared orBOND PAIRS and unshared orLONE PAIRS. This is called a LEWIS ELECTRON DOTstructure.

  25. Rules of the Game •# of valence electrons of an atom = Group # •For Groups 1A-4A, # of bond pairs = group # • For Groups 5A -7A, BP’s = 8 - Group # •Except for H (and sometimes atoms of 3rd and higher periods), BP’s + LP’s = 4 This observation is called the OCTET RULE

  26. Building a Dot Structure Ammonia, NH3 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. C, N, P, & S are common central atoms. Therefore, N is central 2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs

  27. H H N H •• H H N H Building a Dot Structure 3. Form a single bond between the central atom and each surrounding atom 4. Remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair.

  28. Sulfite ion, SO32- Step 1. Central atom = S Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2 TOTAL = 26 e- or 13 pairs Step 3. Form bonds 10 pairs of electrons are now left.

  29. •• O • • • • •• •• O S O • • • • •• •• Sulfite ion, SO32- Remaining pairs become lone pairs, first on outside atoms and then on central atom. •• Each atom is surrounded by an octet of electrons.

  30. Carbon Dioxide, CO2 1. Central atom = carbon 2. Valence electrons = 16 or 8 pairs 3. Form bonds. This leaves 6 pairs. 4. Place lone pairs on outer atoms.

  31. Carbon Dioxide, CO2 4. Place lone pairs on outer atoms. 5. So that C has an octet, we shall form DOUBLE BONDS between C and O.

  32. Double and even triple bonds are commonly observed for C, N, P, O, and S H2CO SO3 C2F4

  33. OR bring in bring in right pair left pair •• •• •• O S O • • • • •• •• Sulfur Dioxide, SO2 1. Central atom = S 2. Valence electrons = 18 or 9 pairs 3. Form double bond so that S has an octet — but note that there are two ways of doing this.

  34. Sulfur Dioxide, SO2 This leads to the following structures. These equivalent structures are called RESONANCE STRUCTURES. The true electronic structure is aHYBRIDof the two.

  35. Urea, (NH2)2CO 1. Number of valence electrons = 24 e- 2. Draw bonds.

  36. Urea, (NH2)2CO 3. Place remaining electron pairs in the molecule.

  37. Urea, (NH2)2CO 4. Complete C atom octet with double bond.

  38. VSEPR (valence-shell electron-pair repulsion) Theory (p.78-80 IB book) • states that the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible. • Unshared pairs of electrons are also important in predicting the shapes of molecules.

  39. Carbon dioxide (CO2) 180° No unshared electron pairs on carbon Linear shape (180º bond angle) • The carbon in a carbon dioxide molecule has no unshared electron pairs. • The double bonds joining the oxygens to the carbon are farthest apart when the O=C=O bond angle is 180°. • Thus, CO2 is a linear molecule.

  40. Planar Triangular Shape (120º bond angle) • CH2O

  41. Tetrahedral shape (109.5º bond angle) • The hydrogens in a methane molecule are at the four corners of a geometric solid called a regular tetrahedron

  42. Structure Determination by VSEPR Ammonia, NH3 The electron pair geometry is tetrahedral. The MOLECULAR GEOMETRY — the positions of the atoms — is PYRAMIDALwith a 107˚ bond angle

  43. Structure Determination by VSEPR Water, H2O The electron pair geometry is TETRAHEDRAL The molecular geometry is BENT with a 105˚ bond angle.

  44. Inter-molecular Forces INTERmolecular forces. Forces between molecules, between ions, or between molecules and ions. 3 types: 1) induced dipole 2) dipole-dipole 3)hydrogen bonding

  45. INDUCED (temporary) DIPOLES How can non-polar molecules such as O2 and I2 dissolve in water? The water dipole INDUCES a dipole in the O2 electric cloud.

  46. Dipole-Dipole (permanent) Forces Dipole-dipole forces bind molecules having permanent dipoles to one another.

  47. Hydrogen Bonding A special form of dipole-dipole attraction, which enhances dipole-dipole attractions. H-bonding is strongest when X and Y are N, O, or F

  48. H-bond H-Bonding Between Methanol and Water - + -

  49. Hydrogen Bonding in H2O H-bonding is especially strong in water because • the O—H bond is very polar • there are 2 lone pairs on the O atom Accounts for many of water’s unique properties. -ice floats -high surface tension -low vapor pressure

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