Chapter 7 — States of Matter and Changes of State

1. Chapter 7 — States of Matter and Changes of State

2. What are the States of Matter? • Solids • Liquids • Gases

3. Solids • Solids have a definite volume and a definite shape • Molecules (or atoms or ions) are closely packed together • INTERMOLECULAR FORCES are strong enough to hold molecules (etc.) rigidly in place with respect to each other

4. Liquids • Liquids have a definite volume but nodefinite shape • Liquids have the ability to flow • Molecules are very close together, but can flow past each other • Intermolecular forces are • Strong enough to hold molecules in a condensed phase • Not strong enough to prevent molecules from sliding past each other

5. Gases • Gases have neither definite shape nor definite volume • Gases are ideally independent molecules • Intermolecular forces are essentially zero between gas molecules

6. Kinetic - Molecular Theory of Matter • Gases are well-described by gas laws • No such laws exist for solids or liquids. • Particles that make up solid and liquid samples are touching; therefore, solids & liquids are not easily compressible (both are called the condensed states of matter). • K M Theory of matter attempts to describe all the states of matter and the conversion between states • To understand the states of matter, we need to consider the molecules comprising matter and how they interact

7. Ionic Bonds (b/w metals and nonmetals) • An ionic bond results from Coulombic attraction between oppositely charged ions (like charges repel each other and opposite charges attract each other). • Ions are not molecules • Ionic compounds are (almost) always solids under normal conditions • Metals give up 1 or 2 electrons to achieve a noble-gas-like electron configuration (cations) • Nonmetals acquire 1 or 2 electrons to achieve a noble-gas-like electron configuration (anions)

8. Covalent Bonds (b/w nonmetals) • Covalent bonds result from sharing one or more pairs of electrons • The OCTET RULE states atoms want to be a Noble gas (filled valence shell with 8 electrons). • Or at least have an noble gas electron configuration • This is achieved be giving each atom access to 8 electrons • Every atom tends to add, remove, or share electrons so as to end up with eight valence electrons. • Valence electrons: electrons in the highest energy (outermost) shell (valence shell). • Species with the same electron configurations are termed isoelectric (every atom “strives” to be isoelectric to the nearest noble gas). • When nonmetals combine, neither one can force its partner to become an ion so each nonmetal atom has access to as many electrons as the nearest noble gas (sharing electrons)

9. e1¯ 2 e2¯ Valence Bond Theory • A covalent bond results from overlap of two electron clouds after bringing atoms close enough • e- cloud b/w 2 nuclei will “shield” them from each other reducing repulsion • This allows the electron on atom 1 to spend time around the nucleus of atom 2 and vice versa (e.g., 2 hydrogen atoms overlapping) Electron on atom 1 can “zip over and bask” in the positive glow of the nucleus on atom 2 and viceversa http://www.chemistryland.com/CHM130W/11-Bonds/bonds.html e1¯ 1 + + e2¯

10. Bond Length and Bond Energy (atoms too close to each other, repulsion b/w positive nuclei, rapid rise of energy, decreased stability (atoms are separated so no covalent bond) (amount of energy needed to brake the bond) (energy is at a minimum and stability at maximum)

11. Lewis Dot Structures • G.N. Lewis developed the theory of covalent bonding • Structures showing covalent bonds are called Lewis structures • Each line represents a shared pair of electrons • Lone pairs of electrons are shown by a pair of dots

12. Drawing Lewis Structures http://www.whfreeman.com/chemicalprinciples/content/instructor/sampleproblems.pdf • Decide on atom connectivity • Hydrogen is frequently bonded to oxygen • Oxygen is rarely the central atom • Count the total number of valence electrons • An atom’s number of valence electrons is equal to its group number • Connect the atoms with single bonds • A single bond is one shared pair of electrons • Use lone pairs and/or multiple bonds to give each atom an octet of electrons

13. How many valence electrons are in the following? • N Nitrogen is in group 5A. It has five valence electrons. • H2S Hydrogen has one valence electron, and sulfur has six. The total for the molecule is 2(1) + 6 = 8. • CO32– Carbon has four valence electrons; oxygen has six; then two for the charge. 4 + 3(6) + 2 = 24. • NH4+ Nitrogen has five valence electrons; hydrogen has one, minus one for the charge. 5 + 4(1) – 1 = 8.

14. Example H2O Atom connectivity H-O-H Total # Valence electrons? 8 (6 electrons from oxygen, group 6A + 1 electron per hydrogen, group 1A) 6 + 2 = 8 (divide by 2 = 4 electron pairs) -use 2 of the pairs to give single bonds b/w the oxygen and each hydrogen. -place two remaining electron pairs on oxygen as lone pairs of electrons. CO2 Atom connectivity O-C-O Total # Valence electrons? 16 (6 electrons per oxygen, group 6A = 12 + 4 electrons from carbon, group 4A) 12 + 4 = 16 (divide by 2 = 8 electron pairs) -use 4 of the pairs to give double bonds b/w the carbon and each oxygen. -place 2 of the remaining electron pairs on one oxygen as lone pairs of electrons and the other 2 pairs on the other oxygen.

15. Example (Isoelectric Species & Triple Bonds) CO CN– Cyanide is CN-, which means that there are 4 shell electrons for carbon, 5 for nitrogen and one more to make the anion. This makes a total of 10 electrons. Nitrogen naturally wants to make three bonds, carbon wants four, but two atoms can't have a quadruple bond between them. So they have three covalent bonds, which uses up 6 electrons. This leaves four electrons, which divides up nicely into two lone pairs (one for each atom). So the lewis dot structure is:

16. VSEPR Theory • Valence Shell Electron Pair Repulsion Theory (VSEPR) • Atoms are bound into molecules by shared pairs of electrons • Electrons repel each other (like charges repel each other) • Therefore, the groups bonded to a central atom try to get as far apart from each other as possible • The goal is to minimize electron pair repulsions around a given central atom

17. Molecular Geometry: describes spatial arrangement of the atoms in a molecule (intermolecular forces are caused by electrons are arranged in a molecule) Linear Geometry • AX2 (“A” is central Atom; “Xs” are “stick-on” groups bonded to central Atom) e.g., CO2 • Bond angle b/w C & O: 180° HCN XeF2

18. Bent • AX2E2 (A is central Atom; Xs are stick-on groups; Es are lone pairs of electrons on the central Atom • Bond angle: ~105° H2O SO2

19. Trigonal Planar Urea (NH2)2CO • AX3 (no lone pairs of electrons) • Bond angle: 120° BF3 CO32—

20. Pyramidal • AX3E (three bonding pairs and one lone pair all point to the vertices of a tetrahedron but we only consider the bonding pairs in molecular geometry) • Bond angle: ~108° NH3

21. Tetrahedral • AX4 (central Atom with 4 stick-on groups) • Bond angle: 109.5° CH4 (methane) CHCL3 (chloroform)

22. Electronegativity • Linus Pauling was the first to develop an electronegativity scale • Electronegativity is the tendency of an atom in a molecule to attract shared electrons to itself.1(Atoms “pulling” of electrons to themselves) • Fluorine is the most electronegative element (EN = 4.0) • The closer an atom is to fluorine, the more electronegative it is (O is more electronegative than N; Cl is more electronegative than Br) 1Chemistry, 7th Edition, Zumdahl & Zumdahl

23. Polar Covalent Bonds • If two atoms of identical electronegativity are bonded together, the bond is non-polar (no “electron hot spots”) • If two atoms of different electronegativity are bonded together, the bond is polar, and the electrons spend more time around the more electronegative atom • This creates partial charges (“electron hot spots”) • The greater the difference in EN between two atoms, the more polar the bond (the bonding electrons spend more time around he more electronegative atom since they are “pulled” to such atom; therefore the more time the “shared” electron pair spends on that atom) • The limiting example of this is the ionic bond (don’t confuse ionic bonds with polar covalent bonds!)

24. Example • The bond in hydrogen is non-polar • The bond in hydrogen chloride is polar • Chlorine is more electronegative than hydrogen so the bonding pair of electrons “spends” more time around Cl; therefore, the Cl end of the molecule has a partial negative charge (delta+) and the hydrogen end of the molecule has a partial positive charge (delta-) (direction in which the bond is polarized)

25. Molecular Polarity • Bond dipoles are vectors • The vectoral sum of the bond dipoles gives the molecular dipole

26. Example • Predict the molecular dipole for water Bond dipole vectors act together to give a net molecular dipole

27. Example • Predict the molecular dipole for carbon dioxide Bond dipole vectors cancel each other out to resulting in a nonpolar molecule (even though the molecule has polar bonds)

28. Intermolecular Forces • These are attractive forces between molecules or atoms or ions • Immensely important • These forces hold DNA molecules in a helix and are the mechanism for DNA transcription • There are three main varieties of IM forces: • Dipolar, hydrogen bonding, and London forces

29. d- d- d+ d+ H H Cl Cl Dipole – Dipole Attraction (aka “dipolar attraction) • This is the attraction between the opposite (partial) charges of polar molecules (dipolar attractive force)

30. London Forces • Also called Van der Waals forces, these are created by instantaneous dipoles • London forces are much weaker than either dipole-dipole or H-bonding (ubiquitous) • London forces get stronger with larger atoms/molecules because larger molecules have more electrons.

31. He He London Forces Between Helium Atoms For the merest fraction of time, there is a dipole-dipole attraction between the atoms.

32. O—H O—H H H Hydrogen Bonding (special type of dipolar interaction) • This is generally stronger than dipolar attractions (strongest intermolecular forces) • Hydrogen bonding is the attraction between a hydrogen bonding directly to an F, O, N (all of these 3 are highly electronegative). Hydrogen bonding is very important intermolecular force (page 166)

33. Ion – Dipole Attraction • This is the attraction between an ionic charge and a polar molecule • This attraction allows ionic solids to dissolve in water • The strength of this force varies widely and depends on the magnitude of the dipole moment of the polar species and the size of the ion

34. A Sodium Ion and a Chloride Ion Hydrated by Water Molecules (very polar) Na Cl

35. Effects of Intermolecular Forces • More intermolecular forces mean: • Higher boiling and melting points • Higher heats of fusion and vaporization • Lower vapor pressure • More viscous liquids • IM Forces also affect solubility • “like dissolves like”

36. Explain the trend in boiling points of the halogens(temperature at which the vapor pressure is equal to the ambient pressure (1 atm)

37. Which of These Will Be Soluble in Water? • HCl(g) (polar) • O2(g) (nonpolar)

38. Which of These Is a Gas? • N2O or NaN3 ? • Cl2 or I2 ?

39. Changes of State

40. Vapor Pressures • The most energetic molecules in a liquid have sufficient kinetic energy to escape into the gas phase • Once the molecules are free as gases, they exert a pressure • This is called the vapor pressure • How does vapor pressure depend on temperature? • Vapor pressure INCREASES with INCREASING temperature

41. Changes of State • Solid to liquid: melting • Liquid to solid: freezing • Liquid to gas: vaporization (evaporation) • Gas to Liquid: condensation • Solid to gas: sublimation • Gas to solid: deposition

42. Energy Changes and Changes of State • Imagine recording the temperature of an 18 gram(i.e.,1.0 mole) sample of ice at -40°C as heat is added

43. Heating Curve for 1 Mole of Water Water is boiling: Heat of vaporization Ice is melting: Heat of fusion

44. Molar Enthalpy of Vaporization(ΔHvap ) • ΔH°vap is the heat required to convert one mole of liquid to a gas at its normal boiling point • ΔH°vap is an inherently endothermic process (amount of energy that most be added to the sample for the phase transition to occur) • ΔH°vap has units of energy/quantity, e.g., kJ/mole • ΔH°vap represents the energy needed to break intermolecular forces and allow molecules to escape into the gas phase

45. Molar Enthalpy of Fusion • ΔH°fus is the heat required to convert one mole of solid to a liquid at its normal melting point • ΔH°fus is an inherently endothermic process • ΔH°fus can have units of kJ/mole • ΔH°fus represents the energy needed to break down intermolecular forces and allow molecules to slide around the liquid phase

46. Question • Why does steam at 100°C cause more severe burns than water at the same temperature? a) When water at 100oC touches your skin, it begins to drop its temperature immediately as the water cools to your skin temperature. b) When steam at 100oC touches your skin, it remains at that temperature while releasing the heat of vaporization onto your skin as the gas converts to a liquid. Then you have water at 100oC as “a” above.

47. Some Heats of Vaporization & Fusion

48. Average KE # Molecules Kinetic Energy Distribution of Energy • In a sample of material, the kinetic energies of the molecules (etc.) follow a Boltzmann Distribution:

49. Dynamic Equilibrium • At the surface of a liquid, the most energetic molecules can escape from the IM forces into the gas phase • Gas molecules near the surface of a liquid can be captured by IM forces into the liquid state • When there is a balance between vaporization and condensation, a state of dynamic equilibrium exists

50. Vapor Pressure • Molecules can escape from the surface of a liquid into the gas phase • The gaseous molecules exert a pressure, call the vapor pressure • The vapor pressure of a liquid increases with increasing temperature • This is quantitatively described by the Claussius-Claperyon equation: