Unit 12 Chapters 11, 12, 13, 25

Organic Molecules Hydrocarbons Compounds containing only carbon and hydrogen Carbon always makes 4 bonds N – 3, O – 2, H – 1 Alkanes are all single bonds Alkenes have a double bond Alkynes have a triple bond

Condensed Structural Formula Shows who is bonded to who, without the actual bonds H H H │ │ │ H─ C─ C─ C─ H H H H │ │ │ H─ C─ C= C─ H │ H

Functional Groups Specific arrangement of atoms that give compounds a unique property Usually involve more than C, H’s Can be on the end of a chain, in the middle, or separating chains

Organic Molecules Halides or halocarbons Alcohols Amine Halogen attached to a carbon Alcohols –OH group attached to Carbon (–O–H) Amine Nitrogen attached to a carbon (–NH2)

Organic Molecules O ǁ ─ C ─ H Aldehyde Carbonyl group at end of chain Name ends with –al Condensed structural formula ends with -CHO H H O │ │ ǁ H─ C─ C─ C─ H │ │ H H H H H H O │ │ │ │ ǁ H ─ C─ C─ C─ C─ C─ H │ │ │ │ H H H H Propanal Pentanal CH3CH2CHO CH3CH2CH2CH2CHO

Organic Molecules O ǁ ─ C ─ O ─ H Organic acid Carboxyl group on end of chain Name ends in –oic acid Condensed structural formula ends with -COOH Hydroxyl H is the acidic H H O │ ǁ H ─ C─ C─ OH │ H Ethanoic Acid CH3COOH

Organic Molecules O H ǁ ǀ ─ C ─ N ─ H Amide ǁ ǀ ─ C ─ N ─ H Amide Carbonyl group with an amine group attached to it Must be on an end Name ends in -amide H H O │ │ ǁ H─ C─ C─ C─ NH2 │ │ H H Propanamide CH3CH2CONH2

Organic Molecules Ketone Ether Double bonded oxygen on a middle carbon Name ends with –one Ether Single oxygen between 2 carbon chains H O H │ ǁ │ H─ C─ C─ C─ H │ │ H H Acetone CH3COCH3 H H H H │ │ │ │ H─ C─ C─O ─ C─ C─ H Diethyl Ether CH3CH2OCH2CH3

Organic Molecules O ǁ ─ C ─ O ─ Ester Carbonyl group with single oxygen between carbon chains Named in two parts 1st Branch off oxygen first as alkyl group 2nd Chain containing Carbonyl group Ending in –oate H O H │ ǁ │ H ─ C─ O ─ C─ C─ H │ │ H H Methyl Ethanoate CH3COOCH3

Isomers Two or more compounds with the same chemical formulas, but different structural formulas and different properties Different Names H H H H H │ │ │ │ │ H ─ C─ C─ C─ C─ C─ H H H CH3 H │ │ │ │ H─ C─ C─ C─ C─ H H H H H C5H12 H CH3 H │ │ │ H─ C─ C─ C─ H Pentane Methyl Butane 2,2-Dimethyl Propane

States of Matter The main difference between the states of matter is difference in distance between particles Condensed phases Solids, liquids Fluids Liquids, gases

States of Matter The state of a substance depends on two main factors Kinetic energy of particles Attraction between particles Intermolecular forces of attraction

Intermolecular Forces Three main forces Dispersion nonpolar Dipole-dipole polar Hydrogen bonds H with N, O, F

Dispersion Forces Attraction between electrons of one atom and protons of another atom Induced dipole All atoms and molecules have dispersion forces

Dispersion Forces

Dispersion Forces Attractive forces increase with increasing number of electrons more polarizability Attractive forces increase with increasing mass

Dispersion Forces More surface areas increases attraction

Dipole - Dipole Attraction between polar molecules This is in addition to dispersion forces

Dipole - Dipole

Hydrogen Bonds A hydrogen bonded to an N, O, F will be attracted to another N, O, F N, O, F are both very small and very electronegative Special case of dipole – dipole HF, NH3, H2O….

Hydrogen Bonds

Intermolecular Forces Ion – Dipole Regents: Molecule Ion Ionic compounds dissolved in polar solvents Discussed later in unit with solutions

Solids State of matter with a definite volume and definite shape Particles are packed tightly together Two types Crystalline Amorphous Non-crystalline

Crystalline Solids Particles are in a highly ordered arrangement

Amorphous Solids No particular order to the arrangement

Network Covalent Solids All atoms are connected in a network of covalent bonds Diamonds, graphite, SiO2 Very hard Very high melting and boiling points Usually not conductive

Network Covalent Solids

Solids Type Bond Type MP, BP Conductivity Soluble in H2O Molecular Covalent Low No Polar only Network Very High Ionic High Liquid, Dissolved Most Metallic Solid, Liquid

Semiconductors Material with electrical conductivity between a metal and an insulator Metal > semiconductor > insulator Usually C, Si, Ge

Semiconductors Doping Addition of impurities p-type n-type One less valence electron (positive) Ga n-type One extra valence electron (negative) As

Liquids State of matter with a definite volume, but takes the shape of its container. Particles are close to each other due to intermolecular forces Particles are able to slide past each other due to kinetic energy

Liquids Intermolecular forces play a large role in determining a number of properties of liquids Surface Tension Viscosity Capillary Action Vapor Pressure

Surface Tension Tendency of a liquid to minimize surface area Stronger intermolecular forces cause higher surface tensions

Viscosity Resistance of a liquid to flow Stronger intermolecular forces cause higher viscosity

Affect of Temperature Increasing temperature decreases surface tension Increasing temperature decreases viscosity Decreasing surface area also decreases viscosity

Capillary Action Ability of a liquid to flow up a thin tube against the pull of gravity

Capillary Action Happens because of two forces working together Cohesion Force that holds liquid molecules together Adhesion Attraction of liquid molecules to container walls

Meniscus Curving of a liquid surface in a thin tube Water Mercury Adhesion > Cohesion Mercury Cohesion > Adhesion

Vapor Pressure Pressure exerted by a vapor in equilibrium with liquid Stronger intermolecular forces cause lower vapor pressure More attraction = less evaporation

Solution Homogeneous Mixture Components can’t be filtered Uniform Throughout Components can’t be filtered Particles aren’t large enough to scatter light Tyndall Effect

Separating Mixtures Evaporation – evaporate away liquid to leave solid Distillation – Separates homogeneous liquid mixtures based on different boiling points Chromatography – separation of substances based on polarity and intermolecular forces

Solution Components Solvent Solute Dissolving medium in mixture Usually water Solute Dissolved particles in solution Solute goes into solvent

Molarity Molarity = Moles of Solute Liters of Solution 1 mol/L = 1 M Often used for solids dissolved into liquids Most common concentration system

Solubility Soluble Insoluble Solute will dissolve in solvent Solute will not dissolve in solvent

Solubility Miscible Immiscible Soluble liquid – liquid mixtures Ammonia is miscible in water Immiscible Insoluble liquids Oil is immiscible in water

“Like Dissolves Like” Polar and ionic substances will dissolve in polar solvents Nonpolar substances will dissolve in nonpolar solvents

“Like Dissolves Like” Acetone is soluble in water because of dipole – dipole interactions Iodine will dissolve in hexane because of dispersion interactions

“Like Dissolves Like” NaCl will dissolve in water because of Ion – Dipole interactions Polar ends of water will be attracted to oppositely charged ions of ionic compounds Solvation

Coulomb’s Law There are stronger ion – dipole interactions when ionic charges are larger Ca2+ > Na+ There are stronger ion – dipole interactions when ions are smaller Li+ > Na+

Spectrophotometry Using light to make measurements in chemistry UV – electronic transitions Visible – solution concentration Infrared – molecular vibration Microwave – molecular rotation

Spectrophotometry Devices Spectrophotometer Measures transmittance over a range of wavelengths Colorimeter Measures absorbance of specific wavelengths

Spectrophotometry Devices Spectrophotometer

Spectrophotometry Devices Colorimeter

Beer’s Law Relates the absorbance of light to concentration A = abc A = Absorbance a = molar absorptivity (unique to substance) b = path length c = concentration

Beer’s Law Calibration curve