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Introduction to Chemistry
- 1. 09/10/13 cottingham Introduction to Inorganic Chemistry A Review of the Basic Regents Chemistry Concepts
- 3. 09/10/1309/10/13 cottinghamcottingham MatterMatter • Anything that has mass and takes up space • Everything you “see” around you is composed of matter
- 4. 09/10/1309/10/13 cottinghamcottingham MassMass • The amount of matter an object has • Measured in grams
- 5. 09/10/1309/10/13 cottinghamcottingham ElementElement • A substance that cannot be broken down chemically into simpler substances • About 92 naturally occurring elements – Example – gold, magnesium, neon, etc. • About 25 elements necessary for life • Six most common elements – Carbon, hydrogen, oxygen, nitrogen . . . calcium and phosphorus – >99% of living things • Trace elements – necessary in only small (i.e. trace) amounts – Example – Iodine • Necessary for proper thyroid function • Goiter caused by iodine deficiency
- 8. 09/10/1309/10/13 cottinghamcottingham AtomAtom • Simplest particle of an element that retains the properties of that element • A given atom is unique to a given element • Atoms composed of smaller particles called subatomic particles
- 9. 09/10/1309/10/13 cottinghamcottingham Compounds vs. MoleculesCompounds vs. Molecules • Compound - A substance composed of two or more elements. – Ex. NaCl – sodium chloride (table salt) • Molecule – A substance composed of two or more atoms, can be the same element. – Ex. O2 – oxygen (diatomic) – However, sodium chloride is also a molecule
- 10. 09/10/1309/10/13 cottinghamcottingham The Periodic Table of ElementsThe Periodic Table of Elements ((in any language, it’s still the samein any language, it’s still the same))
- 12. 09/10/1309/10/13 cottinghamcottingham Important DefinitionsImportant Definitions • Atomic Number – # of protons – Unique for each element • Mass number – Sum of the protons and neutrons • Atomic Mass – Average of masses of all of the isotopes of an element – Usually a decimal number very close to the mass number
- 14. 09/10/1309/10/13 cottinghamcottingham Periodic TablePeriodic Table • Elements arranged according to atomic number starting with hydrogen – (atomic # = 1)
- 15. 09/10/1309/10/13 cottinghamcottingham Periodic TablePeriodic Table • Contains horizontal rows called periods – Ascending atomic # from left to right. • Contains vertical columns referred to as groups. – We are concerned with groups IA-VIIIA – Elements in a group have similar chemical and physical properties
- 16. 09/10/1309/10/13 cottinghamcottingham Group IA – Alkali MetalsGroup IA – Alkali Metals • Strong metallic qualities • Highly reactive – Not found alone in nature • One valence electron • Tendency to lose one e- when reacting.
- 17. 09/10/1309/10/13 cottinghamcottingham Group IIA – Alkaline EarthGroup IIA – Alkaline Earth MetalsMetals • Strong metallic qualities • Very reactive – Not found free in nature • Harder and denser than alkali metals • Two valence e- • Tendency to lose 2 e- during reactions
- 18. 09/10/1309/10/13 cottinghamcottingham Group VIA - ChalcogensGroup VIA - Chalcogens • More varied in properties – Oxygen, Sulfer – nonmetals – Selenium, tellurium – metalloids – Polonium – metal • Very Reactive • Six valence e- • Tendency to gain two electrons
- 19. 09/10/1309/10/13 cottinghamcottingham Group VIIA - HalogensGroup VIIA - Halogens • All nonmetals • Often found in diatomic state (F2) • Very reactive • Seven valence e- • Tendency to gain 1 e- during reactions
- 20. 09/10/1309/10/13 cottinghamcottingham Group VIIIA- Noble GasesGroup VIIIA- Noble Gases • Nonmetals • Eight valence e- • Don’t lose or gain electrons • They are nonreactive (inert)
- 21. 09/10/1309/10/13 cottinghamcottingham Metals vs. NonmetalsMetals vs. Nonmetals
- 22. 09/10/1309/10/13 cottinghamcottingham Metals vs. NonmetalsMetals vs. Nonmetals • Metals – Solid at room temperature – Conduct heat and electricity well – Malleable (sheets) – Ductile (wires) – Lustrous (shiny) – High melting/boiling points • Nonmetals – Opposite of metals • Metalloids – Some qualities of both metals and nonmetals
- 25. 09/10/1309/10/13 cottinghamcottingham Trends in the Periodic TableTrends in the Periodic Table
- 26. 09/10/1309/10/13 cottinghamcottingham Structure of the AtomStructure of the Atom
- 27. 09/10/1309/10/13 cottinghamcottingham Did you know?Did you know? • Atoms are mostly empty space? – If the nucleus of an atom was the size of a golf ball, the nearest electron would be roughly 1 km away! • The nucleus of an atom is extremely dense. – The same size nucleus would have a mass of approximately 2.5 billion tons!
- 29. 09/10/1309/10/13 cottinghamcottingham Some More Things to KnowSome More Things to Know • All atoms of a given element have the same # of protons • All atoms are considered neutral in charge unless designated with a symbol of charge, in which case they are considered an ion; # electrons = # protons except in ions • The # of neutrons is equal to or greater than the number of protons. – mass # - atomic # = # neutrons
- 30. 09/10/1309/10/13 cottinghamcottingham 3 Major Subatomic Particles3 Major Subatomic Particles • Proton – Positive charge – 1 x 10 -24 grams (about 1 dalton) – Located in the nucleus • Neutron – Neutral (no charge) – 1 x 10 -24 grams (about 1 dalton) – Located in the nucleus • Electron – Negative charge – 1/2000 the mass of a proton or neutron – Moving in orbitals around the nucleus at about the speed of light
- 31. 09/10/1309/10/13 cottinghamcottingham Energy Levels (electron shells)Energy Levels (electron shells) • Electrons exist at varying energy levels • The further they are from the nucleus, the more energy they have – Think centripetal force • Electrons tend to occupy the lowest energy level (closest to nucleus) possible • Electrons can be “excited” to higher energy levels for very brief periods – Example: Light energy during photosynthesis
- 32. 09/10/1309/10/13 cottinghamcottingham Excitation of an ElectronExcitation of an Electron
- 33. 09/10/1309/10/13 cottinghamcottingham Electron Configuration and ChemicalElectron Configuration and Chemical PropertiesProperties Why atoms reactWhy atoms react • It’s all about the # of valence electrons! • Valence electron shell is the outermost shell (that contains electrons) • A full valence shell = inert (stable electron configuration) • Anything else = reactive (unstable electron configuration)
- 34. 09/10/1309/10/13 cottinghamcottingham Here’s the DealHere’s the Deal • 1st energy level – Full (stable) with 2 electrons • 2nd energy level – Full (stable) with 8 electrons • 3rd energy level – Full (stable) with 8 electrons
- 36. 09/10/13 cottingham Isotopes • Most elements have at least 2 isotopes, some have several. • Isotopes vary in the # of neutrons only. • Example: Carbon has 3 isotopes – 12 C – stable (6 neutrons) – 13 C – stable (7 neutrons) – 14 C – radioactive (8 neutrons)
- 37. 09/10/1309/10/13 cottinghamcottingham Uses of Radioactive IsotopesUses of Radioactive Isotopes • Dating fossils – Carbon – 14 • Measure half-life (5730 years) • Medical tracers – Iodine – 131 • Various types of sensors can detect radiation.
- 38. 09/10/1309/10/13 cottinghamcottingham Shorthand AbbreviationShorthand Abbreviation • 7 3Li • 16 8O • How many protons, neutrons and electrons do the above examples have?
- 39. 09/10/13 cottingham Answers • Lithium generally has – 3 protons – 3 electrons – 4 neutrons • Oxygen generally has – 8 protons – 8 electrons – 8 neutrons
- 40. 09/10/1309/10/13 cottinghamcottingham Lewis Electron Dot DiagramsLewis Electron Dot Diagrams • Show the electron configuration for only the valence e- for an atom • Steps – Write the symbol of the atom – Make a dot for each valence e- (use the “four sides” of the symbol) – Only one rule – don’t pair up e- until after all four orbitals have one e- each
- 41. 09/10/13 cottingham Examples • Lithium: 1 valence e- • Chlorine: 7 valence e-
- 42. 09/10/1309/10/13 cottinghamcottingham PracticePractice • Draw the Lewis dot diagram for the following atoms (use your periodic table) – Hydrogen – Helium – Beryllium – Carbon – Nitrogen – Oxygen – Fluorine – Rubidium – Iodine
- 43. 09/10/13 cottingham Using LEDD’s to Determine Bonding
- 45. 09/10/1309/10/13 cottinghamcottingham 4 Major Types of Bonds4 Major Types of Bonds • Strongest to weakest – Covalent bonds – Ionic bonds – Hydrogen bonds – van der Waals interactions
- 46. 09/10/1309/10/13 cottinghamcottingham Covalent BondsCovalent Bonds • Strongest • Generally occurs when two nonmetals interact • A pair, or pairs, of e- are shared • Single covalent bond – One pair of e- shared between two atoms – Represented by a single line in structural formula • Double covalent bond – Two pairs of e- shared between two atoms – Represented by a double line in structural formula • Triple covalent bond – Three pairs of e- shared between two atoms – Represented by a triple line in structural formula
- 48. 09/10/13 cottingham Structural vs. Molecular Formulae (single covalent bonds) • Methane • Methane CH4
- 49. 09/10/13 cottingham Double and Triple Covalent Bonds (structural formulae)
- 50. 09/10/1309/10/13 cottinghamcottingham Quick PracticeQuick Practice • React hydrogen with fluorine • React hydrogen with oxygen • React hydrogen with carbon • React carbon with oxygen • React Nitrogen with hydrogen
- 51. 09/10/1309/10/13 cottinghamcottingham Polar vs. Nonpolar Covalent BondsPolar vs. Nonpolar Covalent Bonds • It’s all about electronegativity – Electronegativity • The affinity an atom has for electrons – i.e. How strongly it pulls on both its own e- and the e- of other atoms • All atoms are electronegative, some more than others • Polarity, whether or not a molecule is polar or nonpolar, can have a big effect on the behavior of the molecule.
- 52. 09/10/1309/10/13 cottinghamcottingham Nonpolar Covalent BondsNonpolar Covalent Bonds • Occurs between two atoms of the same electronegativity. • Electrons are shared equally – Both atoms are pulling with the same force • Examples – eneg = electronegativity – O=O (O2) • Same atom – same eneg – C—H • Carbon and hydrogen have the same eneg
- 53. 09/10/1309/10/13 cottinghamcottingham Polar Covalent BondsPolar Covalent Bonds • Occurs between two atoms of differing eneg • Electrons are not shared equally – i.e. e- spend more time around one atom than the other • This creates a slight polarity of charge in the molecule – More eneg atom gains slightly negative charge – Less eneg atom gains a slightly positive charge • Note – oxygen is the big one here • Example – H – O bond – Hydrogen is slightly positive, oxygen slightly negative
- 54. 09/10/13 cottingham Water Molecule (polar)
- 55. 09/10/13 cottingham Water Molecule (polar)
- 56. 09/10/1309/10/13 cottinghamcottingham Electronegativity for ImportantElectronegativity for Important AtomsAtoms • F – most eneg • O – highly eneg • N – eneg • Cl – eneg • C and H are middle of the road eneg • C – H bond is nonpolar
- 58. 09/10/1309/10/13 cottinghamcottingham Ionic BondsIonic Bonds • Also strong – Relatively weak around water • Around water, ionically bonded substances dissociate into ions • Generally occur between a metal and a nonmetal – Metal loses electron, nonmetal gains electron • Electrons are not shared, they are transferred from one atom to another • Differences in eneg are great • Ions (charged particles) are formed • An ionic bond is an attraction between oppositely charged ions.
- 59. 09/10/1309/10/13 cottinghamcottingham Examples of Ionic BondsExamples of Ionic Bonds • Na + Cl Na+ + Cl- NaCl – Cl steals an e- from Na, gains a 1- charge and leaves Na with a 1+ charge. The oppositely charged ions are attracted. • Mg + 2F Mg2+ + 2F- MgF2 – Two fluorinessteal 1 e- each from Mg, gain a 1- charge and leave Mg with a 2+ charge. The oppositely charged ions are attracted.
- 60. 09/10/1309/10/13 cottinghamcottingham Trends for Ionic BondingTrends for Ionic Bonding • Group IA, 1+ ions – Except H • Group IIA, 2+ ions • Group VIIA, 1- ions • Group VIA, 2- ions
- 61. 09/10/1309/10/13 cottinghamcottingham PracticePractice • Using LEDD’s . . . – React potassium with iodine – React calcium with chlorine • Answers – KI – CaCl2
- 62. 09/10/1309/10/13 cottinghamcottingham Hydrogen BondsHydrogen Bonds • Hbonding is an attraction between the slightly positively charged atom in one polar bond and the slightly negatively charged atom in a different polar bond • Occur only between polar molecules or polar regions of molecules • Weak, short-lived bonds (still very important) • This can happen between two different molecules or between different regions of the same molecule
- 63. 09/10/13 cottingham Example of a Molecule With Polar Regions (phospholipid)
- 64. 09/10/13 cottingham Hydrogen Bonding Between Two Water Molecules
- 65. 09/10/13 cottingham Hydrogen Bonding Between Regions of the Same Molecule
- 66. 09/10/13 cottingham Hydrogen Bonding in DNA
- 67. 09/10/1309/10/13 cottinghamcottingham Van der Waals InteractionsVan der Waals Interactions • Due to the random movement of electrons • Weak • Short-lived • Can occur in both polar and nonpolar molecules • Only occur when molecules are very close together • Allows all molecules to be attracted to one another • Plays role in the shape of larger molecules
- 68. 09/10/13 cottingham Van der Waals
- 70. 09/10/1309/10/13 cottinghamcottingham Molecular ShapeMolecular Shape • Every covalently bonded molecule has a characteristic size and shape. FOR IB BIO…the only thing about shape to remember is: • Biological Structure is related to function – i.e. A molecule’s structure is directly related to its “job” • Molecules communicate via shape
- 73. 09/10/13 cottingham Cell Surface Receptors
- 74. 09/10/13 cottingham Enzymes (catalyze reactions) http://ntri.tamuk.edu/cell/an-enzyme.gif
- 76. 09/10/13 cottingham Saccharine (Sweet ‘n Low)
- 79. 09/10/1309/10/13 cottinghamcottingham 6CO6CO22 + 6H+ 6H22OO CC66HH1212OO66 + 60+ 6022 • Represented by chemical equations – Reactants on the left – Products on the right – Some bonds are broken and reformed – Mass is conserved in a reaction • In a balanced chemical equation, the total # of atoms of each element must be equal on both sides of the equation • Is this equation balanced?
- 80. 09/10/1309/10/13 cottinghamcottingham EquilibriumEquilibrium • In some reactions, all of the reactants are converted to products • Most reactions, however, are reversible – they can go in either direction • CO2 + H2O H2CO3 • Eventually, equilibrium will be met. – This is when the reaction is occurring in both directions at the same rate
- 81. 09/10/1309/10/13 cottinghamcottingham Activation EnergyActivation Energy • The energy necessary to start a reaction • Can be high • This is good – control • Enzymes (usually proteins) act as catalysts to lower the EA – control
- 82. 09/10/1309/10/13 cottinghamcottingham Exergonic/Endergonic ReactionsExergonic/Endergonic Reactions and Free Energyand Free Energy • Free energy – energy that can be used to do work • Exergonic reactions – release free energy – result in products with less stored energy than the reactants – Reactants (high E) products (lower E) + E (free) – C6H12O6 + 602 6CO2 + 6H2O + E – Molecules are being broken down (catabolism) • Endergonic reactions – store free energy – result in products with more stored energy then the reactants – Reactants (lower E) + E (free) products (high E) – 6CO2 + 6H2O + E C6H12O6 + 602 – Molecules are being built up (anabolism)
- 83. 09/10/1309/10/13 cottinghamcottingham Oxidation – Reduction ReactionsOxidation – Reduction Reactions REDOXREDOX • LEO the lion goes GER • Loses e- oxidation, gains e- reduction • Any time an ion is formed – redox reaction • Example – Na + Cl Na+ + Cl- • Na has lost e- and has been oxidized • Cl has gained an e- and has been reduced
- 84. 09/10/1309/10/13 cottinghamcottingham Redox in Covalent bondsRedox in Covalent bonds • Redox rxns can also involve covalent bonding • Atom can be reduced if it becomes bonded to a highly eneg atom. – i.e. it’s own e- are being pulled away from it • Example – C-H bond broken, H replaced with O, C-O – Oxygen is highly e-neg – Carbon has been oxidized – Oxygen has been reduced
- 85. 09/10/1309/10/13 cottinghamcottingham Dalton’s Atomic TheoryDalton’s Atomic Theory • We already have discussed this, but to make it more clear the following 5 ideas are the keys to make the Atomic Theory more easy to identify • 1. Elements are made of tiny particles called atoms. • 2. All atoms of a given element are identical
- 86. 09/10/1309/10/13 cottinghamcottingham Dalton’s Atomic TheoryDalton’s Atomic Theory • 3. The atoms of a given element are different from those of any other element. • 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms.
- 87. 09/10/1309/10/13 cottinghamcottingham Dalton’s Atomic TheoryDalton’s Atomic Theory • 5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.
- 89. 09/10/1309/10/13 cottinghamcottingham Describing SolutionsDescribing Solutions • A solution is a uniform mixture • Two types of parts – Solvent –the dissolving agent • Water is a great example (especially in cells) – Solutes – are dissolved in the solvent • Anything dissolved in a substance • There can be many solutes in a given solvent • Example – mix salt and water – Water is the _____ – Salt is the _______
- 90. 09/10/1309/10/13 cottinghamcottingham Like Dissolves LikeLike Dissolves Like • Polar vs. nonpolar • Polar and nonpolar substances repel one another • So . . . • Polar (and ionic) solutes will dissolve in polar solvents • Nonpolar solutes will dissolve in nonpolar solvents • Think oil (nonpolar) and vinegar (polar)
- 91. 09/10/1309/10/13 cottinghamcottingham Hydrophobic vs. HydrophilicHydrophobic vs. Hydrophilic • HYDROPHOBIC • Hydro = water • Phobic = fearing • Don’t dissolve in water • Nonpolar substances • OIL • HYDROPHILIC • Hydro = water • Philic = loving • Do dissolve in water • Polar/ionic substances • VINEGAR
- 92. 09/10/1309/10/13 cottinghamcottingham ReviewReview • Like dissolves like • Hydrophilic and hydrophobic, i.e. nonpolar and polar molecules, literally repel one another • All polar molecules are hydrophilic • All ionic molecules are hydrophilic • All nonpolar molecules are hydrophobic • However – Some molecules can be both hydrophobic and hydrophilic (in different areas) – Example – phospholipids
- 93. 09/10/1309/10/13 cottinghamcottingham PHOSPHOLIPIDPHOSPHOLIPID hydrophobic and hydrophilic regionshydrophobic and hydrophilic regions
- 94. 09/10/1309/10/13 cottinghamcottingham Cell MembraneCell Membrane phospholipid bilayerphospholipid bilayer
- 95. 09/10/1309/10/13 cottinghamcottingham Cell MembraneCell Membrane phospholipid bilayerphospholipid bilayer
- 96. 09/10/1309/10/13 cottinghamcottingham Concentration of a SolutionConcentration of a Solution • A measure of the amount of solute/solvent • Lots of solute and/or low solvent = a high concentration (represented by [x] ) • Aqueous solution – water is the solvent – Very important to life • Saturated solution – cannot dissolve any more solute
- 98. 09/10/1309/10/13 cottinghamcottingham Ionic Substance DissolvingIonic Substance Dissolving
- 99. 09/10/1309/10/13 cottinghamcottingham Covalent Substance DissolvingCovalent Substance Dissolving
- 100. 09/10/1309/10/13 cottinghamcottingham Acids and BasesAcids and Bases
- 101. 09/10/1309/10/13 cottinghamcottingham Dissociation into IonsDissociation into Ions • To break into separate ions in solution • Ionically bonded substances do this – NaCl Na+ (aq) + Cl- (aq) • Covalently bonded substances don’t dissociate into ions, with one exception • Water is the “exception” – H2O H+ + OH- • Note H+ and H3O+ are synonymous
- 102. 09/10/1309/10/13 cottinghamcottingham Acids and BasesAcids and Bases • Acids • H3O+ ↔ H+ + H2O • H3O+ = Hydronium • Acidity or alkalinity (bases) is actually a measure of hydronium and hydroxide ions dissolved in a solution • BASES • OH- = hydroxide ion • REMEMBER: NaOH ↔ Na+ + OH-
- 103. 09/10/1309/10/13 cottinghamcottingham AcidsAcids • Proton donors • Increase H+ (proton) concentration • Can be strong or weak • Example of a strong acid – Hydrochloric acid (HCl) – HCl H+ + Cl- – Dissociates completely
- 104. 09/10/1309/10/13 cottinghamcottingham Bases (alkaline)Bases (alkaline) • Proton acceptors • Decrease H+ (proton) concentration • Can be strong or weak • Example of a strong base – Sodium hydroxide (NaOH) – NaOH Na+ + OH- – Dissociates completely – Makes lots of hydroxides which “eat up” protons – OH- + H+ H2O
- 106. 09/10/1309/10/13 cottinghamcottingham pH and lifepH and life • Control of pH is very important to living things (homeostasis) • Example – Human blood pH range generally 7.35 – 7.45 – Anything below 7 or above 7.8 can be deadly • Buffers – Weak acid/base that can neutralize small amounts of another acid/base – H2CO3 H+ + HCO3 - – Carbonic acid hydrogen ion + bicarbonate ion