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Chapter 11

1 st shell. d subshell. 2 nd shell. p. s. d. f. Chemical Bonds: Formation of Compounds. Chapter 11.

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Chapter 11

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  1. 1st shell d subshell 2nd shell p s d f Chemical Bonds: Formation of Compounds Chapter 11 Periods in the Periodic System of elements are called shells; subshells (s, p, d, f…) contain orbitals (s:1, p: 3, d: 5, f: 7), and each orbital can be occupied by 2 electrons with opposite spins. Elements are classified in groups having similar chemical characteristics. Noble gases have completely filled p orbital with 8 e-. 4s orbital is populated before 3d. Before we continue with filling 4p, 10 e- must go into 3d. We put them in a loop since we are stalled filling 4th shell. This presentation has 8 groups.

  2. Atomic properties Ionization energy … is the energy required to remove an e- from the atom. and increases bottom to top, up a group. It increases left to right across a period… That is the behavior exactly opposite to that of atomic radii! Ionization Energy vs. Atomic Number Atomic radii increase right to left across the period, and top to bottom down the group.

  3. 74 pm . . H H Chemical Bonding A molecule is a collection of atoms bound together. It is considered as an element if all atoms are of the same type (e.g. H2), or a compound if it is made of different atoms. A bond between two atoms of hydrogen will occur spontaneously if the atoms are within a close proximity, i.e. when attractive forces between the nucleus of one atom and the e- of the other overcome repulsive forces among their e-, thus pulling the atoms together. The bond that is formed is called covalent bond. (Co-means partner, valent refers to valence electrons). It always releases energy. In a covalent bond between two hydrogen atoms, each atom apparently has 2 electrons in its shell. That gives each H atom the electronic configuration of the closest noble gas, He. 1 pm = 10-12 m Filled-shell e- (or core e-) are almost never involved in the bond as they are too close to their own nucleus. Electrons are shared between two atoms. 2 He atoms repel each other and will never form a bond. Covalent bond can be presented as any of the following ways: H : H Remember that the line stands for a pair of e-! H – H

  4. Molecules and the Octet Rule Elements ‘want’ to have the electronic configuration identical to that of noble gas (8e-). When form a molecule, atoms achieve the octet (8e-) by sharing e- with other atoms. Hydrogen is an exception, as it only needs one more electron to fill its 1s orbital. How many electrons an element needs to satisfy the octet rule can be found by observing the Roman numerals in the periodic table. C: 8e- H: 2e- All atoms in the molecule have the electronic configuration identical to that of the closest noble gas. Core electrons Valence electrons O: 8e- H: 2e- Nonmetals usually gain electrons, metals loose them.

  5. . . . . . . . . . . . . . . . . : C – O : : C – O : : O – C – O : : O – C – O : : O – C – O : . . . . . . . . . . . . . . . . O = C = O . . . . : O – C Ξ O : . . : O Ξ C – O : . . Step 1: Find the total number of valence e- by adding up the group numbers of all atoms. For ions, adjust the dot count accordingly (subtract e- for cation, add for anion). Drawing Dot Diagrams CO2 CO Group IVA (14) Group IVA Group VIA Group VIA (16) + 2 x 6e- 4e- + 4 e- 6 e- Step 2: Unless told otherwise, assume that the first non-hydrogen atom in the formula of the group is the central atom. Connect atoms with single bonds. Total: 16e- Total: 10e- O – C – O C – O 16e – 4e = 12e- in lone pairs The central atom must form multiple bonds, hence H can never be the central atom. or Step 3: Subtract 2e- for each bond from total #e- to get#e- in lone pairs. 10e – 2e = 8e- in lone pairs Step 4: Put in the remaining electrons, two at a time, as lone pairs. Start with the terminal atoms, and continue with the central atom if there are any electron pairs left. or . . . . Step 5: Check that each atom has octet satisfied (doublet for H). If not, move electron pair(s) from the adjacent atom to form multiple bonds. . . : C Ξ O : . . Practice with CO32-, SO3, etc. Equivalent structures, or resonance forms.

  6. 2- : O O : C . . . . O : : . . C2H4 C2H2Cl2 CO32- Which C is the central atom? 2x4e + 2x1e + 2x7e = 24e- 4e + (3 x 6e) + 2e = 24e- total 2x4e+4x1e=12e-total BOTH ! total H H O O C C cannot pull electrons from Cl 12e – (5x2e) = 2e-inlone pairs H H C C O C C 24e – (3 x 2e) = 18e- : Cl : : Cl : H H . . . . in lone pairs H H C C . . : . . : : Cl H Cl H : : . . . . H H H H : O O : C C C C C . . . . C C Octet rule not satisfied for C O : : . . : : : H Cl : Cl H . . . . H H Octet rule not satisfied for C H has no e- to give in for double bonds . . Two more resonance structures for CO32- ion are possible. HW chapter 11: 1, 3, 13, 29, 37 SO3 is identical to this; try it yourself.

  7. The bond between metals and non-metals is usually ionic. Metals give away their e- and become positively charged (cations). Nonmetals accept them and become anions. The ionic bond is formed as a result of attraction between oppositely charged ions. The compound is called ionic compound, and the three-dimensional ordered network of the ions is called ionic lattice. Electrons are rarely shared equally between atoms. Electronegativity and the Polar Covalent Bond Electronegativity (EN) is numerical rating of an atoms ability to attract to itself the shared electrons in a covalent bond. Generally, electronegativity of metals is low, and that of nonmetals is high. The least electronegative atom (except H!) is the central atom in dot structures. Polar covalent bond is a covalent bond in which e- are shared unequally (large DEN). A partial negative charge (d-) occurs on the more EN atom. A partial positive charge (d+) occurs on the less EN atom. The difference in EN defines the bond. DEN = 0, covalent; DEN = 1.0, polar covalent (23% ionic); DEN = 1.9, polar covalent (60% ionic); DEN > 1.9, ionic. Ionic and covalent are two extremes at the ends of a continuum bonding types.

  8. : O : | | C H H The Shape of the Molecules H | H – C – H | H Valence Shell Electron Pair Repulsion (VSEPR) theory is the model mostly used to predict molecular shape. Electron pairs on the central atom repel one another. The two dimensional dot structure of methane, CH4. gives the angles between electron pairs of 90o. But the dot structure angles are arbitrary. Molecules are three dimensional, and the electron pairs would be further away if the third dimension is considered. In fact, the shape of methane molecule is tetrahedral; the bond angles between electron pairs is 109.5o. Four electron pairs around an atom assume tetrahedral arrangement. When there are not enough electrons for single bonds the molecule forms multiple bonds and the structure differs. VSEPR theory treats each multiple bond as a single electron group, because it occupies roughly the same region of space. The number of electron groups around an atom is called the atom’s steric number (SN). Dot structures of formaldehyde and acetylene are arbitrarily shown with angles of 90o. : O : | | H – C – H H H | | C Ξ C H – C Ξ C – H Their true geometry has bond angles of 120o and 180o, respectively. formaldehyde acetylene Dot structures True geometry

  9. If the central atom has a lone pair of electrons, that electron pair is included in the molecular shape. VSEPR arrangements of electron groups around an atom having no lone pair electrons The dot diagram of ammonia presents the atom in a plane. O = C = O The steric number on nitrogen is 4 (3 bonding pairs and a lone pair). The e- pairs on the N assume tetrahedral arrangement. . . H – N – H | H Electrons in the lone pair occupy more space than the bonding pairs. They squeeze H-N-H bond angle to approx. 107o. We describe the shape of the molecule, not that of its electrons. Lone pair(s) of electrons are thereforeignored. The molecule of NH3 has a pyramidal shape. Rule of thumb: each lone pair of e- on a period-2 atom compresses the remaining bond angles around that atom by ~2o.

  10. H | O O | | O H O Using VSEPR 1. Draw a dot diagram. 2. Count the number of e- pairs around the central atom, including lone pairs (i.e. the steric number, SN). A multiple bond counts as a single e- group. 3. Find the best arrangement of the electrons using SN. 4. Pretending the lone e- pairs are invisible, describe the resulting shape of the molecule. Practice on H2O and O3. H-O-H bond angle ~105o. bent : : : O-O-O bond angle 120o. SN=4 SN=3 bent

  11. evaporation liquid gas condensation Water and Properties of Liquids Chapter 13 Liquids have intermediate properties between solids and gases. Liquids are almost incompressible, have definite volume and assume the shape of the container. Densities of liquids are usually lower than that of their solids. Water is an exception. Evaporationorvaporization is the escape of molecules from liquid into gaseous state. During evaporation, liquid that stays behind is cooler. The opposite process is condensation. Sublimation is the escape of molecules directly from solid into gas, bypassing liquid state. Open container completely evaporates. Vapor pressure is the pressure exerted by a gas at equilibrium with its liquid, so that: Closed container reaches equilibrium between liquid and gas. Vapor pressure depends only on temperature, not on the amount of liquid.

  12. Vapor Pressure Measurement 1 atm = 760 torr 20 oC 20 oC 30 oC 20 oC a. c. b. d. a. The system is evacuated. c. Equilibrium established. Manometer attached to the flask shows equal pressure in both legs. Manometer shows constant pressure difference, 17.5 torr. d. Temperature raised to 30 oC. b. Water is added. Equilibrium reestablished. Manometer shows constant pressure difference of 31.8 torr. Liquid evaporates. Manometer shows increase in pressure.

  13. Vapor pressure and temperature Vapor pressure of any gas at the boiling point is equal to the atmospheric pressure. Vapor pressure of ethyl ether is the highest at any temp. TBP TBP Vapor pressure: Ether > Alc. > Water. TBP Rate of evaporation: Ether > Alc. > Water. proportional to vapor pressure. Volatility Boiling point: Substances that readily evaporate are volatile. Ether < Alc. < Water Volatile Vapor pressure of ethyl ether at 20 oC: 442.2 torr Moderately volatile Vapor pressure of water at 20 oC: 17.5 torr Vapor pressure of mercury at 20 oC: 0.0012 torr Nonvolatile

  14. Boiling Point Curves 1 atmosphere pressure evaporation liquid gas TBP ethyl ether 34.6oC TBP alcohol 78.4oC condensation melting solid liquid freezing Normal Boiling Point Boiling point at standard pressure (1 atm, or 760 torr). Each point on the curve represents a vapor-liquid equilibrium at a particular temperature and pressure. At 500 torr, ethyl ether boils at ~22 oC, alcohol at ~68 oC, and water at 89 oC. TBP water 100.0oC Freezing or Melting Point The temperature at which the solid and liquid are in equilibrium. Changes of State Heating curve for a pure substance Majority of substances change phases upon heating: solid  liquid  gas. CO2 is an exception (dry ice sublimes). A – B: solid state B – C: melting C – D: liquid state D – E: evaporation E – F: vapor state Temperature is constant during melting and boiling – all heat used to break solid (at boiling point) or liquid forces.

  15. Heat of Fusion and Heat of Vaporization We learned before that amount of heat depends on mass and temp. change. Qheating = (mass) (spec.heat) (temp.change) Energy (heat) needed to change 1 g of a solid at its melting point into liquid is heat of fusion. Energy (heat) needed to change 1 g of a liquid at its boiling point into vapor is heat of vaporization. Constant temperature! Qfusion = (mass) (spec.heat of fusion) Qvaporization = (mass) (spec.heat of vaporization) Example 1: How many joules is needed to change 20.0 g of ice at 0 oC to steam at 100. oC? Qheating = (mass) (spec.heat) (temp.change) Qtot = Qfusion + Qheating + Qvaporization } Qfusion = (20.0 g) x (335 J/g) Qtot = 60.3 kJ Qheating = (20.0 g) x (4.184 J/goC) x (100. oC) Qvaporization = (20.0 g) x (2260 J/g) Hydrogen Bond produces unusually high melting & boiling point

  16. . . . . H H H – O : H – O : | . . | | | H – C – O – C - H H H | . . | H H Hydrogen Bonding (cont.) H bonding exists between H directly bonded to one of the three most electronegative elements (Fluorine, Oxygen, and Nitrogen), and F, O or N of another molecule. H bond No H bond . . . H bonds are intermolecular forces. No H bonded to F, O, or N H bonded to O Ethyl ether Surface Tension and Capillary Action A droplet of liquid falling forms a sphere due to attractions to other liquid molecules – surface tension. Cohesive forces within mercury liquid (left) are stronger than adhesive forces between Hg and walls of the container. Opposite is true for H2O. Spontaneous rise of liquid in a narrow tube – capillary action.

  17. Hydrates . . . . O H H Some ionic solutions retain water upon evaporation. It becomes the part of the crystalline compound – water of crystallization. The formula is written as: ionic compound, dot , # water molecules… . CuSO4 5 H2O and name them by adding # (Latin) hydrate. Copper(II) sulfatepentahydrate. dry CuSO4 – white Hydrate = blue Hydrates are true compounds and the water is an integral part of it. Formula mass CuSO4.5 H2O: 63.55+32.07+64.00+5x18.02 = 249.7 Percent composition of water is (5x18.02 / 249.7) x 100 = 36.08% Water indicator CuSO4. 5 H2O(s)  CuSO4(s) + 5 H2O(g) Water can be removed by intense heat: The reaction is reversed when water is added. d- Water, a Unique Liquid Water covers ~75% of Earth. 97% of water is in the oceans. Only 3% is fresh water, of which 2/3 is locked up in ice polar caps. d+ Solid form (ice) has lower density than liquid water. Water is very stable molecule, can stand temperatures up to 2000 oC. It does not conduct electricity when pure, but decomposes into H2 and O2 in solutions of ions. 2 H2 + O2 --> 2 H2O + 484 kJ Water can be formed by Combustion, Neutralization, Metabolic reaction 2 C2H2(g) + 5 O2 4 CO2 + 2 H2O(l) + 1212 kJ HCl(aq) + NaOH(aq) --> NaCl(aq) + 2 H2O C6H12O6(aq) + 6 O2 6 CO2(g) + 6 H2O(l) + 2519 kJ

  18. Reactions of water Water reactions with metals: Cold water reacts with Na, K, Ca: Na + H2O  H2 + NaOH Steam reacts with Zn, Al and Fe: Fe + H2O(g) --> H2 + Fe3O4 Remind yourself of the activity series: the above six metals are the most active. Another three metals are more active than H: Pb, Sn, and Ni and react with acids only; Cu, Ag, Hg and Au are below H in the series and do not react with acids or H2O. Water also reacts with certain nonmetals. Anhydride means: without water. Most reactive: 2 F2 + 2 H2O(l) --> 4 HF(aq) + O2 Less reactive: Cl2(g) + H2O(l)  HCl(aq) + HOCl(aq) To test whether a metal or nonmetal is an anhydride, try to remove H2O until all hydrogen is removed. Least reactive: C(s) + H2O(g)  CO(g) + H2(g) Water reacts with metal and non-metal oxides: Basic anhydride: CaO + H2O(l)  Ca(OH)2(aq) D OH Ca CaO + H2O Acid anhydride: CO2(g) + H2O(l)  H2CO3(l) OH D Water Purification H2SO4 SO3 + H2O Screening, flocculation and sedimentation, sand filtration, aeration, disinfection. Hard water contains Mg2+ and Ca2+ ions Additional water purification is done by distillation, Ca2+, Mg2+ precipitation, ion exchange and demineralization. HW (p.332): 7, 11, 17, 25

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